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A  TEXT-BOOK 


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INORGANIC    CHEMISTRY. 


BY 
DR.    A.    F.    HOLLEMAN, 

Professor  Ordinarins  in  the  University  of  Amsterdam;  Emeritus  Profess&i 

Ordinarius  in  the  University  of  Groningen,  Netherlands,  and 

Fellow  of  the  Royal  Academy  of  Sciences,  Amsterdam. 


ISSUED  IN  ENGLISH  IN  COOPERATION  WITH 
HERMON   CHARLES   COOPER. 


FOURTH   ENGLISH  EDITION,   COMPLETELY  MWISED. 
TOTAL    ISSUE,    SIXTEEK   THOUSAND. 


NEW  YORK 

JOHN  WILEY  &  SONS, 

LONDON:     CHAPMAN    &     HALL, 
1913 


x 


Copyright,  1901,  1902,  1905,  1908,  1911, 

BY 

HERMON  C.  COOPER. 


First,  Second,  and  Third  Editions  entered  at  Stationers'  Hall. 


lUE  SCIENTIFIC   PRESS 

ROBERT   DRUMMOND  AND   COMPANY 

BROOKLYN,   N.   Y. 


PKEFACE  TO  THE  FOUETH  EDITION. 


THE  present  edition  represents  a  thorough  revision  of  the  work 
by  the  Dutch  author  and  the  American  collaborator.  It  profits 
by  the  author's  experience  with  the  frequent  editions  in  other 
languages  but  is  independent  in  composition. 

Very  many  of  the  descriptive  portions  have  been  rewritten, 
notably  those  on  the  sulphur  oxides  and  acids,  rare  gases,  nitrogen 
oxides  and  acids,  sodium  hydroxide  and  carbonate,  radio-active 
elements  and  platinum,  as  well  as  the  sections  on  thermo- 
chemistry, colloids  and  the  iron-carbon  system,  while  the  sub- 
jects of  stability  and  the  reality  of  molecules  and  atoms  furnish 
new  material.  The  chapter  on  metal-ammonia  compounds  is 
reprinted  as  approved  by  Professor  WERNER  for  the  third  edition. 

Notwithstanding  the  appearance  of  differential  formulae  in 
the  book,  it  is  believed  that  a  student  who  is  unfamiliar  with 
the  calculus  should  have  little  difficulty  in  understanding  the 
meaning  and  use  of  such  formulas,  provided  he  is  willing  to  take 
the  author's  word  for  the  solutions  of  the  equations. 

Independent  students  may  well  be  cautioned  against  regard- 
ing any  text-book  as  infallible.  Even  in  a  book  with  a  world 
market,  such  as  this  one  enjoys,  undergoing  many  revisions  by 
the  author  and  by  collaborators  in  other  nations,  and  being 
frequently  reviewed  critically  by  the  journals,  there  will,  probably, 
always  be  some  textual  errors  and  some  passages  whose  lucidity 
could  be  improved.  Readers  can  therefore  render  great  service 
by  reporting  all  unsatisfactory  passages  to  t|j|  publishers. 

Thanks  are  herewith  expressed  to  my  colleague,  Professor  H. 
MONMOUTH  SMITH,  for  constant  aid  in  detect  ing:  errors. 

References  in  the  text  to  "  ORG.  CHEM."  refer  to  the  companion 
volume  of  this  work,  HOLLEMAN'S  "  Text-book  of  Organic  Chem- 
istry," translated  by  WALKER  and  MOTT. 

H.  C.  COOPER. 

SYRACUSE  UNIVERSITY, 
October,  1911. 


CONTENTS. 

\ 

Light-face  figures  refer  to  pages;  heavy -face  figures  to  paragraphs. 

PAGE 

INTRODUCTION  (1-5) 1 

PHYSICAL  AND  CHEMICAL  PHENOMENA  (6) 3 

CHEMICAL  OPERATIONS  (7) 5 

THE  ELEMENTS  (8) 7 

Oxygen  (9,  10) 9 

Law  of  HENRY,  11;  Oxidation,  12;  Analytic  and  synthetic  methods, 
13. 

Hydrogen  (11-13) 13 

Oxyhydrogen  blowpipe,  15;  Detonating-gas,  15;  Reduction,  16. 

THE  CONSERVATION  OF  MATTER  (14) 16 

Water  (15-19) 17 

Physical  properties,  20;  Natural  water,  20;  Composition  of  water,  22. 

COMPOUNDS  AND  MIXTURES  (20) 25 

PHENOMENA  ACCOMPANYING  THE  FORMATION  OR  DECOMPOSITION 

OF  A  COMPOUND,  27. 
EXPLANATION  OF  THE  CONSTANT  COMPOSITION  OF  COMPOUNDS;  ATOMIC 

THEORY  (21-23) 27 

Law  of  constant  composition,  27;  Atoms,  28;  Molecules,  28;  Law 
of  multiple  proportions,  29;  THE  ATOMIC  WEIGHTS  OF  THE 
ELEMENTS,  29;  CHEMICAL  SYMBOLS  AND  FORMULAE,  30. 

STOICHIOMETRICAL  CALCULATIONS  (24) .  .  . 31 

Chlorine  (25-35) 33 

Catalytic  action,  34;  Hydrogen  chloride,  37;  Acids,  bases  and  salts, 
39;  Composition  of  hydrochloric  acid,  41;  Law  of  GAY-LUSSAC, 
43;  AVOGADRO'S  hypothesis,  44;  Molecular  weights,  45;  RULES 
FOR  DETERMINING  MOLECULAR  AND  ATOMIC  WEIGHTS,  47; 
Kinetic  theory,  47;  General  gas  equation,  48;  THE  REALITY  OF 
MOLECULES  AND  ATOMS  AND  THEIR  ABSOLUTE  WEIGHT,  49. 

Ozone  (36,  37) $. 50 

Formula  of,  52;  Allotropism,  53. 

Hydrogen  Peroxide  (38,  39) .* 53 

Status  nascendi,  54. 

MOLECULAR  WEIGHT  FROM  THE  MEASUREMENT  OF  THE  DEPRESSION  OF 
THE  FREEZING-POINT  AND  ELEVATION  OF  THE  BOILING-POINT  (40-43) .     57 
Semi-permeable  membranes,  57;    Osmotic  pressure,  58;    PFEFFER'S 
experiments,  59;    Isotonic  solutions,  62;    Formula  of  hydrogen 
peroxide,  65. 

v 


vi  CONTENTS. 


Bromine  (44,  45) 65 

Hydrogen  bromide,  67. 

Iodine  (46-48) 69 

Hydrogen  iodide,  71. 

DISSOCIATION  (49-51) 72 

Reversible  reactions,  73;  Equilibrium,  73;  Reaction  velocity,  75; 
Law  of  chemical  mass  action,  75;  Unimolecular  and  bimolecular 
reactions,  76. 

Fluorine  (52,  53) 79 

Hydrogen  fluoride,  81. 

Compounds  of  the  halogens  (54-62):    with  each  other,  83;    with 

oxygen,  83. 

NOMENCLATURE  (63),  92. 
Summary  of  the  halogen  group  (64),  93.  . 

ELECTROLYTIC  DISSOCIATION  (65,  66) 94 

Ionic  equilibrium,  98;  Strength  of  acids  and  bases,  99;  Hydrolysis, 
101. 

Sulphur  (67-93) 102 

Allotropic  modifications,  104;  THE  TRANSITION  POINT,  106;  "STA- 
BLE," "METASTABLE,"  AND  "LABILE,"  108;  THE  PHASE  RULE 
OF  GIBBS,  109;  Hydrogen  sulphide,  115;  Solubility  product,  119; 
Hydrogen  persulphide,  120;  Compounds  of  sulphur  with  the  hal- 
ogens, 121;  VALENCE,  122;  Compounds  of  sulphur  with  oxygen, 
124;  Oxygen  acids  of  sulphur,  131;  Volumetric  analysis,  146. 

Selenium  and  Tellurium  (94,  95) 148 

Selenium,  148;  Tellurium,  150. 
Summary  of  the  oxygen  group,  (96)  151. 

THERMOCHEMISTRY  (97-104) 152 

Law  of  HESS,  153:  CHEMICAL  AFFINITY,  156;  THE  DISPLACEMENT 
OF  EQUILIBRIUM,  160;  PASSIVE  RESISTANCES,  161. 

Nitrogen  (105-130) 162 

The  atmosphere,  165;  Argon,  helium  and  companion  elements,  170; 
Compounds  of  nitrogen  and  hydrogen,  174;  Compounds  with  the 
halogens,  179;  Hydroxylamine,  180;  Compounds  with  oxygen, 
181;  Oxygen  acids,  187;  Derivatives  of  the  nitrogen  acids,  195; 
Other  nitrogen  compounds,  198. 

Phosphorus  (131-154) .  .« 199 

Hydrogen  compounds,  204;  Halogen  compounds,  209;  Oxygen  com- 
pounds, 211;  Acids,  212. 

Arsenic  (155-164) 221 

Hydrogen  arsenide,  223;  Halogen  compounds,  226;  Oxygen  com- 
pounds, 226;  Oxy-acids,  228;  Sulphur  compounds,  229;  Sulpho- 
salts,  230. 

Antimony  (165-169) 231 

Hydrogen  antimonide,  232;  Halogen  compounds,  233;  Oxygen  com- 
pounds, 234;  Sulphur  compounds,  236. 


CONTENTS.  vii 

PAGE 

Bismuth  (170-174) 236 

Summary  of  the  nitrogen  group  (175),  239. 

Carbon  (176-189) 241 

Allotropic  forms,  241;  Molecular  and  atomic  weight,  246;  Com- 
pounds with  hydrogen,  248;  Compounds  with  oxygen,  249;  Other 
carbon  compounds,  255;  The  flame,  256. 

Silicon  (190-196) 261 

Hydrogen  silicide,  262;  Halogen  compounds,  263;  Oxygen  com- 
pounds, 265;  Silicic  acids,  266;  COLLOIDS,  268. 

Germanium  (197) 273 

Tin  (198-202) 274 

Stannous  compounds,  276;   Stannic  compounds,  279. 

Lead  (203-206) 281 

Oxides,  283;  Halogen  compounds,  285;  Other  lead  salts,  286;  Sum- 
mary of  the  carbon  group  (207),  287. 

METHODS  OF  DETERMINING  ATOMIC  WEIGHTS  (208-212) 288 

Law  of  DULONG  and  PETIT,  289;  Law  of  NEUMANN,  291;  Law  of 
MITSCHERLICH,  292;  EXPERIMENTAL  DETERMINATION  OF  EQUIVA- 
LENT WEIGHTS,  293. 

THE  PERIODIC  SYSTEM  OF  THE  ELEMENTS  (213-221) 296 

Construction  of  a  system  of  the  elements,  303;  Ascertaining  atomic 
weights,  305;  Prediction  of  properties  of  elements,  307;  Correct- 
ing atomic  weights,  308;  Graphic  representation,  308. 

Lithium  and  Sodium  (222-226) 310 

Lithium,  310;  Sodium,  311;  Oxides  and  hydroxides  of,  312;  Salts 
of,  315;  (Soda,  319). 

Potassium  (227-231) 322 

Oxygen  compounds,  323;  Salts,  324. 

Rubidium  and  Ccesium  (232) '. 329 

Summary  of  the  group  of  alkali  metals  (233),  330. 

Ammonium  Salts  (234) 331 

SALT  SOLUTIONS  (235-239) 335 

ACIDIMETRY  AND  ALKALIMETRY  (240,  241) 350 

Indicators,  352. 

Copper  (242-244) 353 

Cuprous  compounds,  355;  Cupric  compounds,  358. 

Silver  (245-247) ' 359 

Compounds,  362;  Photography,  364. 

Gold  (248-251) 367 

Testing  of  gold  and  silver,  369;  Aurous  compounds,  369;  Auric  com- 
pounds, 370;  Summary  of  the  group  (252),  371. 

Beryllium  and  Magnesium  (253-255) 372 

Beryllium,  372;  Magnesium,  374;  Magnesium  salts,  375. 

Calcium,  Strontium,  and  Barium  (256-262) 376 

Calcium,  376;  Oxides  and  hydroxides  of,  377;  Salts  of,  379;   Glass, 


viii  CONTENTS. 


384;  Strontium,  386;  Barium,  387;  Summary  of  the  group  of  the 
alkaline  earths,  388. 

SPECTROSCOPY  (263-265) 389 

THE  UNITY  OF  MATTER  (266) 395 

Radio-active  Elements  (267) 397 

Zinc  (268,  269) 407 

Cadmium  (270) 410 

Mercury  (271-274) 410 

Amalgams,  412;  Mercurous  compounds,  413;  Mercuric  compounds, 
414;  Summary  of  the  zinc  group  (275),  417. 

ELECTROCHEMISTRY  (276-281) 418 

Boron  (282,  283) 432 

Halogen  compounds,  433;  Oxygen  compounds,  433. 

Aluminium  (284-287) 436 

Compounds  of,  437. 

Gallium,  Indium,  and  Thallium  (287,  288),  441. 

Summary  of  the  group  (289),  442. 

The  Rare  Earths  (290) 443 

Titanium,  Zirconium,  and  Thorium  (291) 446 

Vanadium,  Niobium,  and  Tantalum  (292) 448 

Chromium  Group  (293-299) 449 

Chromium,  449;  Chromous  compounds,  450;  Chromic  compounds, 
450;  Chromates,  452;  Molybdenum,  455;  Tungsten,  456;  Ura- 
nium, 457;  Summary  of  the  group,  458. 

Manganese  (300,  301) 458 

Manganic  acid  and  permanganic  acid,  460. 

Iron  (302-308) 463 

Iron-carbon  system,  466;  Ferrous  compounds,  473;  Ferric  com- 
pounds, 474. 

Cobalt  and  Nickel  (309-312) 477 

Cobalt,  477;  Nickel,  479. 

Platinum  Metals  (313-316) 481 

Ruthenium,   482;    Osmium,   483;     Rhodium,   483;     Iridium,   484; 

Palladium,  484;  Platinum,  485. 

Metal-ammonia  Compounds.      WERNER'S  EXTENSIONS  OF  THE  NOTION 
OF  VALENCE  (317-318) .486 


INORGANIC  CHEMISTRY. 


INTRODUCTION. 

1.  Chemistry  is  a  branch  of  the  natural  sciences, — the  sciences 
which  deal  with  the  things  on  the  earth  and  in  the  outside  universe. 
The  knowledge  of  these  things  is  obtained  by  observation  with 
our  senses,  this  being  the  only  means  we  possess.     It  is  well  to 
understand,  therefore,  that  we  know  not  the  things  themselves, 
but  simply  the  impressions  which  they  make  upon  our  sense-organs. 
When  we  see  an  object,  we  perceive,  in  reality,  only  the  effect  on 
our  retina;  if  we  feel  the  object,  it  is  not  the  body  itself  but  the 
excitement  of  the  sensory  nerves  of  touch  in  our  fingers  that  we  are 
made  aware  of.    Hence  it  may  be  fairly  asked  whether  the  objects 
of  which  we  are  cognizant  are  really  just  as  we  perceive  them,  or 
whether  they  even  exist  at  all  outside  of  our  person.    The  natural 
sciences  leave  this  problem  out  of  consideration — its  solution  is  the 
task  of  speculative  philosophy.     In  reality  they  are  not  concerned 
with  the  objects,  which  in  themselves  we  cannot  know,  but  with 
the  study  of  the  sensations  that  we  receive.    The  sensations  take 
the  place  of  the  objects,  and  we  regard  them  as  such. 

2.  The  Scientific  Investigation   of  Things. — What   is   to   be 
understood  by  the  term?    In  the  first  place,  a  most  accurate 
description  of  the  objects.     From  a  study  of  this  it  is  found  that 
many  objects  resemble  each  other  to  a  greater  or  less  degree,  and 
it  is  therefore  possible  to  make  a  classification,  i.e.  an  arrangement 
of  like  objects  into  groups  and  a  separation  of  the  various  groups 
from  each  other.     By  the  descriptive  method  we  are  finally  able 
to   divide  the  natural  sciences  into  Zoology,  Botany,  Mineralogy 
and  Astronomy. 

3.  In  the  second   place,  scientific  investigation  includes  the 


2  INORG AttlC  CHEMISTRY.  [§  §  3- 

study  of  the  relations  which  the  objects  bear  to  each  other;  in 
other  words,  the  study  of  phenomena.  The  heavenly  bodies  move 
towards  each  other;  water  turns  to  ice  on  cooling;  wood  burns 
when  heated.  It  is  the  task  of  the  natural  sciences  to  accurately 
observe  and  describe  such  phenomena,  i.e.  to  ascertain  in  what 
way  the  heavenly  bodies  change  their  relative  positions,  what 
conditions  affect  the  freezing  of  water,  what  becomes  of  the  burn- 
ing wood,  under  what  conditions  it  can  burn,  etc. 

The  description  of  the  phenomena  leads  to  a  different  division 
of  the  natural  sciences  than  the  description  of  objects,  viz.,  a  divi- 
sion into  Physics,  Chemistry  and  Biology,  the  latter  being  the  study 
of  vital  processes,  and  including  Physiology,  Pathology  and  Thera- 
peutics. 

4.  The  human  mind,  in  pursuing  the  scientific  study  of  nature, 
does  not  feel  contented  with  the  accurate  description  of  objects 
and  phenomena;  it  seeks  also  for  an  explanation  of  the  latter. 
The  various  attempts  at  explanations  constitute  the  most  im- 
portant part  of  science.  When,  for  instance,  we  see  that  a  ray  of 
light  in  passing  through  a  piece  of  Iceland  spar  is  split  up  into  two 
other  rays  of  different  properties,  we  strive  to  account  for  the 
phenomenon.  When  copper  is  heated  in  the  air,  it  turns  into  a 
black  powder;  the  question  again  arises,  why  this  thing  is  so.  In, 
searching  for  an  explanation  of  the  phenomena  we  thus  endeavor 
to  penetrate  deeper  into  the  essence  of  things  than  is  possible  by 
direct  observation.  Although  the  phenomena  themselves  are 
found  to  be  unchangeable,  our  explanation  of  them  may  be  modi-' 
fied  as  our  knowledge  increases.  The  transformation  of  copper 
into  a  black  powder  on  heating  in  the  air  was  formerly  explained 
by  the  supposition  that  something  left  the  metal;  subsequently, 
when  the  phenomenon  was  better  understood,  by  assuming  that 
the  copper  takes  up  something  from  the  air. 

Scientific  investigation  pursues  in  general,  then,  the  following 
course:  A  phenomenon  is  observed  and  studied  as  carefully  as 
possible.  Thereupon  an  explanation  of  it  is  sought.  A  hypoth- 
esis is  set  up.  From  this  conclusions  can  be  formed,  some  of 
which  can  be  tested  by  experiment.  If  the  latter  really  leads  to 
the  expected  result,  the  hypothesis  gains  in  probability.  If  it  is 
subsequently  found  to  explain  and  link  together  a  whole  serie0  of 
phenomena,  it  becomes  a  theory. 


6.]  PHYSICAL  AND  CHEMICAL  PHENOMENA. 

The  nineteenth  century  was  an  era  of  great  prosperity  for 
scientific  inquiry.  For  numerous  phenomena  explanations  have 
been  found  which  possess  a  great  degree  of  probability.  Still  it 
cannot  be  denied  that  the  present  theories  penetrate  only  a  little 
into  the  real  essence  of  things,  and  the  investigator  very  soon 
stumbles  upon  questions  whose  explanation  does  not  at  present 
even  seem  to  be  a  possibility.  The  chemical  process  that  goes  on 
when  copper — to  retain  our  former  example — is  heated  in  the 
air  is  well  known.  However,  the  deeper  question,  why  the  action 
takes  place  just  so  and  not  otherwise,  or  why  the  resulting 
powder  is  black,  still  awaits  a  satisfactory  answer. 

5.  We  observed  in  the  preceding  paragraph  that  the  natural 
phenomena  are  found  to  be  unchangeable.    The  movement  of  the 
planets,  for  example,  still  takes  place  in  the  same  manner  as  in  the 
times  of  the  Ptolemies;    whenever  water  turns  to  ice  the  same 
increase  of  volume  is  to  be  observed;  the  crystal  form  of  common 
salt,  whenever  and  wherever  examined,  is  invariably  the  same; 
from  the  burning  of  wood  the  same  products  are  always  obtained ; 
the  microscopic  structure  of  the  leaves  of  one  and  the  same  plant 
is  never  found  to  vary.    This  general  principle  finds  its  expression 
in  the  phrase,  constancy  of  natural  phenomena.    Every  one  is  con- 
vinced of  its  truth,  and  it  is  tacitly  accepted  as  the  basis  of  every 
natural  scientific  investigation.    If,  for  example,  one  has  measured 
the  angles  which  the  faces  of  a  soda  crystal  form  with  each  other, 
he  considers  it  certain  that  all  soda  crystals  must  show  the  same 
angles,  at  whatever  time  or  place  they  may  be  measured.    If  it  has 
once  been  determined  that  pure  alcohol  boils  at  78°  under  normal 
pressure,  it  is  forthwith  assumed  that  this  must  be  the  case  with 
all  alcohol,  no  matter  how  it  may  be  obtained  or  when  and  where 
it  may  be  tested. 

PHYSICAL  AND  CHEMICAL  PHENOMENA. 

6.  It  was  stated  above  (§3)  that  the  description  of  phenomena 
leads  to  a  division  of  the  natural  sciences  into  Physics,  Chemistry 
and  the  study  of  vital  processes  (Biology) .     In  defining  the  province 
of  Chemistry  Biology  may  be  left  out  of  consideration;   however, 
it  is  desirable  to  compare    the  field  of    Chemistry  with  that  of 
Physics.    In  general  it  may  be  said  that  Physics  deals  with  the 


4  INORGANIC  CHEMISTRY.  [§§  6- 

temporary,  Chemistry  with  the  lasting,  changes  of  matter.  By 
matter  or  substance  we  understand  the  objects  without  reference 
to  their  form.  Iron,  marble,  sand  and  glass  are  kinds  of  matter, 
or  substances,  independent  of  their  external  shape. 

A  couple  of  illustrations  may  make  this  conception  of  temporary 
and  lasting  changes  clear. 

(a)  A  platinum  wire  glows  when  held  in  a  colorless  gas-flame. 
On  removal  it  cools  off  and  no  change  is  visible.  This  is  a  physical 
phenomenon;  the  change,  the  glowing,  is  of  a  temporary  sort.  So 
soon  as  the  cause  of  the  change  is  removed,  the  wire  resumes 
its  original  condition.  When  some  magnesium  wire  is  held  in  the 
flame,  it  burns  with  the  emission  of  a  brilliant  light  and  turns  into 
a  white  powder,  which  is  wholly  different  from  the  substance 
magnesium.  Here  a  lasting  change  has  occurred;  we  have  to  do 
with  a  chemical  phenomenon. 

(6)  Again,  we  may  take  two  white  crystallized  substances,  naph- 
thalene and  cane-sugar,  and  heat  each  separately  in  a  retort  with 
receiver.    The  naphthalene  at  first  melts;  on  continued  heating  it 
begins  to  boil,  then  distils  over  and  condenses  in  the  receiver.    The 
distilled  naphthalene  resembles  the  undistilled  in  every  respect.    The 
substance  has,  as  a  result  of  heating,  undergone  physical  changes — 
melting,  change  to  vapor  and,  finally,  return  to  the  solid  state.    The 
cane-sugar  behaves  differently.    Here  also  a  melting  is  observed  at 
first,  but  soon  the  sugar  turns  darker;    a  brownish  liquid  distils 
over;  a  peculiar  odor  is  noticeable  and  at  last  there  remains  in  the 
retort  a  charred,  porous  mass.    The  cane-sugar  suffers  a  lasting 
change  on  being  heated.     In  this  case  we  have  a  chemical  change, 
(c)  As  a  third  and  last  example  we  may  consider  the  behavior 
of  a  metallic  wire  on  the  one  hand  and  that  of  acidulated  water 
on  the  other,  when  an  electric  current  passes  through  them.    The 
wire  displays  other  properties  so  long  as  the  current  is  on.     If 
the  latter  ceases,  the  wire  returns  to  its  original  condition.    This 
is  a  physical  action.     In  the  acidulated  water,  however,  the  current 
induces  an  evolution  of  gas,  and  this  gas  arising  from  the  water 
has  properties  entirely  different  from  those  of  the  water.    A  lasting 
change  in  the  substance  has  occurred;  a  chemical  action  has  taken 
place, 

A  sharp  distinction  between  physical  and  chemical  phenomena 
is  often — as  will  be  seen  later — very  difficult  to  make. 


7.J  CHEMICAL  OPERATIONS.  5 

CHEMICAL  OPERATIONS. 

7.  In  order  to  avoid  repetitions  it  seems  advisable  at  this  point 
to  describe  briefly  some  of  the  commonest  chemical  operations. 

Solution. — When  sugar,  salt  or  saltpetre,  for  example,  is  put 
into  water,  the  solid  substance  disappears  and  its  taste  is  taken 
on  by  the  water.  The  substance  has  dissolved  in  the  water.  There 
is  a  definite  limit  to  the  solubility  of  each  of  these,  for,  if  the  tem- 
perature is  kept  constant  and  more  of  the  substance  is  gradually 
added,  a  point  is  finally  reached  when  the  water  will  take  up  no 
more.  The  solution  is  then  saturated.  The  solubility  of  most 
solids  increases  with  the  temperature.  Moreover  it  is  very  differ- 
ent with  different  substances,  varying  all  the  way  from  solubility  in 
all  proportions  to  imperceptible  solubility.  Thus  cane-sugar  is  dis- 
solved in  large  quantity  by  water,  while  sand  is  practically  insolu- 
ble in  it.  Liquids  can  be  either  miscible  in  all  proportions 
(water  and  alcohol)  or  only  partially  soluble  in  each  other.  When, 
for  instance,  water  is  shaken  with  a  sufficient  quantity  of  ether 
and  allowed  to  stand,  two  liquid  layers  are  formed;  the  water  has 
dissolved  some  ether  and  the  ether  some  water.  In  most  cases 
the  solubility  of  liquids  in  each  other  also  increases  with  the  tem- 
perature. In  the  case  of  gases  solubility  decreases  with  rising 
temperature. 

Separation  of  a  Solid  and  a  Liquid. — This  may  be  accomplished 
by  filtration.  A  funnel  is  lined  on  the  interior  with  "filter-paper  " 
and  the  mixture  poured  upon  it.  The  solid  is  retained  on  the 
paper  while  the  liquid  passes  through.  Decantation  is  a  less  com- 
plete method  of  separation,  since  more  liquid  remains  with  the 
solid  by  this  method  than  by  filtration.  However,  it  is  evident 
that  neither  method  affords  a  really  complete  separation.  This 
can  only  be  accomplished  by  washing,  i.e.  by  replacing  the  por- 
tion of  the  liquid  which  remains  between  the  solid  particles  by 
another  liquid.  If  the  liquid  of  the  mixture  be  a  salt  solution, 
pure  water  is  very  effective.  It  is  obvious  that  by  repeating  the 
washing  several  times  the  salt  solution  can  be  wholly  removed. 
Suppose  that  1  c.c.  salt  solution  remains  between  the  particles  of 
the  solid  and  that  9  c.c.  water  is  then  added.  The  solution  is  thus 
reduced  to  one-tenth  of  its  original  concentration.  If  1  c.c. 
of  this  dilute  solution  again  remains  with  the  solid  and  another 


6  INORGANIC  CHEMISTRY.  [§§  7- 

9  c.c.  water  is  added,  the  concentration  is  then  10"~2,  or  one- 
hundredth  of  the  original  concentration;  after  six!  such  operations 
it  would  be  only  10~6,  or  one  millionth  of  the  original,  so  that  the 
separation  is  practically  complete. 

Crystallization. — If  a  solution  is  saturated  in  the  warm  and 
'is  then  allowed  to  cool,  the  dissolved  substance  frequently  separates 
out  in  the  crystallized  state.  Advantage  is  often  taken  of  this 
for  purifying  crystallizable  substances. 

Distillation  (Fig.  1). — This  operation  is  frequently  made  use 
of  in  working  with  liquids.  The  liquid  is  placed  in  a  flask  or  a 
retort  and  heated  to  boiling.  The  escaping  vapor  is  cooled  to 


FIG.  1. — DISTILLATION. 

a  liquid  in  a  condenser.  The  latter  consists  of  a  sufficiently  wide 
tube  encased  in  a  jacket,  through  which  water  flows  to  keep  the 
inner  tube  cold.  The  condensed  liquid  is  collected  in  the  receiver. 
It  is  readily  seen  how  volatile  substances  can  be  separated  from 
non-volatile  ones  by  distillation,  e.g.  water  from  salt,  since  the 
former  distil  over  and  the  latter  remain  in  the  distilling-flask. 
However,  liquids  of  different  volatility  can  also  be  separated  in 
this  manner.  Take,  for  example,  a  mixture  of  alcohol  and  water. 
The  more  volatile  constituent,  alcohol,  passes  over  for  the  most 
part  in  the  early  stage  of  the  operation;  towards  the  end  the  less 
volatile,  water.  If  the  two  distillates  are  collected  separately,  an 
approximate  separation  results.  A  few  repetitions  of  this  so-called 
fractional  distillation  bring  about  a  practically  complete  separation 
in  many  cases. 


8.]  THE  ELEMENTS. 

Sublimation. — Certain  solids,  e.g.  camphor,  when  heated  (at 
ordinary  pressure),  turn  to  vapor  without  melting.  If  this  vapor 
comes  in  contact  with  a  cold  surface,  the  substance  is  deposited  in 
the  solid,  crystallized  state.  It  is  evident  that  we  have  here  another 
method  of  separating  some  substances. 

THE   ELEMENTS. 

8.  When  a  substance  (§  6)  is  subjected  to  various  influences, 
such  as  heat,  electricity,  or  light,  or  is  brought  in  contact  with 
other  substances,  it  is  very  often  split  up  into  two  or  more  dis- 
similar components.  As  an  example  let  us  take  gunpowder. 
Water  is  added  and  the  whole  is  stirred  well  and  gently  warmed; 
after  a  while  it  is  filtered,  and  that  which  remains  on  the  filter  is 
found  to  be  no  longer  gunpowder,  for  it  is  unexplosive.  On  evapo- 
rating the  filtrate  a  white  crystalline  substance,  saltpetre,  remains. 
The  undissolved  part  is  dried  and  then  shaken  with  another  sol- 
vent, carbon  disulphide.  After  a  time  the  mixture  is  filtered,  as 
before,  and  there  is  left  on  the  filter  a  black  mass,  consisting  of 
charcoal  powder.  The  carbon  disulphide  of  the  filtrate  evaporates 
and  leaves  yellow  crystals  of  sulphur.  Thus  we  see  that,  by  suc- 
cessive treatment  with  water  and  carbon  disulphide,  gunpowder 
can  be  separated  into  three  substances,  viz.  carbon,  sulphur  and 
saltpetre.  The  two  former  are  incapable,  even  when  subjected 
to  all  the  agencies  at  our  command,  of  division  into  different  com- 
ponents. Not  so  with  saltpetre,  for  when  the  latter  is  heated 
strongly  a  gas  is  given  off  in  which  a  glowing  wooden  splinter  is 
at  once  ignited.  When  the  evolution  of  gas  ceases,  a  substance 
remains  which  gives  off  red  fumes  on  treatment  with  sulphuric  acid, 
something  that  saltpetre  does  not  do.  Saltpetre  can  evidently 
be  broken  up  still  farther  by  heating. 

If  we  subject  all  sorts  of  substances  to  a  successive  treatment 
with  reagents  of  the  most  different  kinds,  we  finally  discover  cer- 
tain ones  that  cannot  be  resolved  into  simpler  substances  by  our 
present  means.  Such  substances  are  called  elements.  Although 
the  number  of  substances,  according  to  §  6,  may  be  considered  as 
infinitely  great,  experience  has  taught  that  the  number  of  elements 
is  small.  There  are  about  eighty. 

As  our  methods  of  examination  improve,  it  may   quite  possibly 


8  INORGANIC  CHEMISTRY.  [§  8 

be  found  that  the  substances  which  the  chemist  of  to-day  regards  as 
elements  have  no  right  to  the  name.  Therefore,  when  we  use  the 
word  "element/'  it  is  to  be  regarded  as  a  relative  term,  dependent 
on  the  extent  of  our  knowledge  and  the  means'  at  our  command.  In 
the  history  of  chemistry  some  cases  are  to  be  found  where  sub- 
stances, once  believed  to  be  elements,  were  subsequently  decom- 
posed. 

The  exact  number  of  elements  cannot  be  definitely  stated,  be- 
cause, on  the  one  hand,  not  all  the  substances  that  possibly  exist 
may  be  within  our  reach,*  and,  on  the  other  hand,  it  is  doubtful 
whether  certain  substances  now  regarded  as  elements  cannot  be 
divided  by  means  already  known. 

On  the  inside  of  the  back  cover  will  be  found  a  list  of  the 
elements  now  known. 

As  may  be  seen  from  this  list,  the  metals  are  included  in  the 
elements.  Together  with  them  we  find  a  number  of  other  sub- 
stances, as  oxygen,  sulphur,  phosphorus,  etc.,  that  are  classed  under 
the  term  non-metals,  or  metalloids.  To  the  latter  class  belong 
many  very  important  substances,  e.g.,  oxygen,  an  element  that 
combines  with  almost  all  others,  causing  what  is  called  combustion. 
Oxygen  is  present  in  a  large  amount  in  the  air.  Another  non- 
metal  is  carbon,  which  is  present  in  all  organized  substances,  and 
is  therefore  a  constituent  of  every  animal  and  plant.  Sulphur, 
which  burns  with  a  blue  flame,  giving  off  a  pungent  odor,  and 
chlorine,  a  greenish-yellow  gas  of  very  disagreeable  odor,  which 
combines  readily  with  most  metals,  are  also  non-metals. 

The  elements  occur  in  very  unequal  proportions  in  the  part  of 
the  earth  accessible  to  us.  Oxygen,  which  occurs  in  air,  in  water, 
and  in  the  solid  part  of  the  earth's  crust,  is  very  preponderant, 
composing  approximately  50%  of  these  portions  of  the  earth  which 
have  been  investigated.  The  elements  silicon,  aluminium,  iron, 
calcium,  carbon,  magnesium,  sodium,  potassium,  and  hydrogen, 
together  with  oxygen,  make  up  99%  of  the  earth's  crust.  There 
remains,  therefore,  only  1%  for  all  the  other  elements.  Some  of 
these  are  quite  common,  e.g.,  lithium,  but  they  almost  always. 

*  Of  the  interior  of  the  earth  only  a  very  small  part  is  known.  If  we  think 
of  the  earth  as  about  the  size  of  an  orange,  the  deepest  mine-shafts  would 
not  even  penetrate  the  thin  yellow  exterior  layer  of  the  orange  skin. 


§  9.]  OXYGEN.  9 

occur  in  very  small  quantities.  Others,  like  niobium  and  tantalum, 
are  found  in  relatively  very  small  amounts  and  in  isolated  places, 
With  the  aid  of  spectroscopy  (§§  263-265),  it  has  been  ascer- 
tained that  the  heavenly  bodies  contain  most  of  the  elements 
found  in  our  earth,  and  also  some  others. 

OXYGEN. 

i 

9.  Under  ordinary  conditions  of  temperature  and  pressure, 
oxygen  is  a  colorless  and  odorless  gas,  whose  most  noticeable 
property  is  its  ability  to  set  glowing  substances  on  fire  with  the 
evolution  of  much  light  and  heat.  A  glowing  splinter  of  wood,  for 
example,  when  introduced  into  an  atmosphere  of  oxygen,  begins 
at  once  to  burn  brightly.  This  action  is  ordinarily  used  as  a 
characteristic  test  for  the  identification  of  oxygen. 

This  gas  can  be  obtained  in  various  ways.  There  are  many 
substances  which  are  known  to  evolve  oxygen  on  heating. 

(1)  Mercuric  oxide,  when  heated  strongly  in  a  retort  (Fig.  2), 
yields  oxygen,  which  can  be  collected  by  means  of  a  delivery-tube 


FIG.  2. — PREPARATION  OF  OXYGEN  FROM  POTASSIUM  CHLORATE. 

opening  under  the  mouth  of  a  cylindrical  receiver  filled  with 
water.  The  inside  of  the  retort  becomes  covered  with  drops  of 
mercury. 

(2)  The  same  apparatus  can  be  used  in  making  oxygen  from 
potassium  chlorate  (chlorate  of  potash),  as  well  as  from  potassium 
nitrate  (saltpetre),  potassium  permanganate,  and  many  other  sub- 
stances. The  preparation  of  oxygen  by  heating  potassium  chlorate 
is  a  method  frequently  used  in  the  laboratory. 


10  INORGANIC   CHEMISTRY.  [§  9- 

Some  substances  give  off  oxygen  when  heated  together  with 
others,  as  in  the  following  cases : 

(3)  Potassium  dichromate  or  manganese  dioxide,  when  heated 
with  sulphuric  acid ; 

(4)  Zinc  oxide,  when  heated  in  a  current  of  chlorine. 

The  atmospheric  air  consists  principally  of  oxygen  and  nitrogen. 
The  following  method  for  separating  these  gases  was  employed  by 
LAVOISIER  in  1774.  He  introduced  some  mercury  into  a  retort  A 
(Fig.  3)  with  a  long,  doubly-bent  neck  that  opened  under  a  bell- 
jar  P  filled  with  air  and  resting  in  a  dish  R  of  mercury.  He  then 


FIG.  3. — ABSORPTION  OF  OXYGEN  BY  MERCURY. 

heated  the  retort  steadily  for  several  days,  keeping  the  mercury 
almost  boiling.  As  a  result,  a  part  of  the  air  in  P  disappeared, 
and  the  gas  remaining  was  found  to  possess  other  properties  than 
air — it  was  nitrogen.  At  the  same  time  the  mercury  had  been 
partially  transformed  into  a  red  powder,  mercuric  oxide.  On 
heating  the  latter  more  strongly  oxygen  was  obtained. 

Oxygen  in  now  prepared  from  liquid  air  (cf.  §  109) . 

The  physical  properties  of  oxygen,  besides  those  already  men- 
tioned, are  as  follows:  Its  specific  gravity,  assuming  the  density 
of  air  to  be  1,  is  1.10535.  A  liter  of  oxygen  at  0°  and  760  mm. 
Hg  pressure  weighs  1.4290  g.  Oxygen  can  be  liquefied;  the 
difficulties  in  obtaining  it  on  a  large  scale  in  the  liquid  state 
have  -now  been  completely  overcome.  Apparatuses  for  lique- 
fying oxygen  have  been  constructed  by  HAMPSON  and  by 
LINDE,  a  description  of  which  is  to  be  found  in  text-books 
of  physics.  The  critical  temperature  of  oxygen  is  -118°, 
and  its  critical  pressure  50  atmospheres.  Liquid  oxygen 


10.]  OXYGEN.  11 

has  a  specific  gravity  of  1.124  (based  on  water)  and  a 
boiling-point  of  -  182.95°  at  745.0  mm.  pressure.  Its  color  is 
light  blue.  It  can  be  preserved  for  some  time  at  ordinary  pres- 
sure, with  the  aid  of  a  so-called  vacuum-flask  (Fig.  4.) 
The  latter  is  a  vessel  enclosed  in  an  air-tight  jacket,  the  space  be- 
tween the  walls  being  evacuated.  100  1.  water  at  0°  dissolves 
4.89  1.  oxygen.  The  gas  is  also  somewhat  soluble  in  alcohol 
and  in  molten  silver.  When  the  silver  solidifies, 
the  oxygen — a  volume  about  ten  times  that  of 
the  metal — suddenly  escapes  from  solution,  caus- 
ing peculiar  elevations  on  the  surface  of  the  silver 
("spitting"  of  silver). 

We  remarked  above  (§  7)  that  the  solubility  of 
gases  in  liquids  diminishes  with  increasing  tem- 
perature. A  very  remarkable  law  expresses  the 
relation  that  exists  between  the  solubility  of  a  gas 
and  its  pressure,  namely,  the  solubility  is  propor- 
tional to  the  pressure.  This  is  the  law  of  HENRY. 
FIG.  4. — VACUUM-  Thus,  when  the  pressure  becomes  a-fold,  the  solu- 
FLASK.  bility  also  becomes  a-fold.  As  the  mass  of  a  gas 
which  is  present  in  a  certain  volume  is  likewise  proportional  to  the 
pressure,  the  law  of  HENRY  can  also  be  expressed  thus :  The  volume 
of  a  gas  dissolving  in  a  certain  quantity  of  a  liquid  is  independent 
of  the  pressure. 

This  law  is  rigid  when  the  solubility  of  the  gas  is  small;  when 
the  solubility  is  large,  for  instance  100  volumes  in  1  of  the  liquid, 
its  deviations  are  considerable. 

Still  another  formulation  of  this  law  is  of  value  in  understanding 
certain  of  its  applications:  The  concentrations  of  the  dissolved  and 
undissolved  portions  of  a  gas  bear  a  constant  ratio  to  each  other.  By 
"  concentration  "  is  meant  the  quantity  of  the  gas  in  grams  per 
unit  volume  (cubic  centimeter). 

10.  Among  the  chemical  properties  of  oxygen  the  most  promi- 
nent is  its  vigorous  support  of  combustion.  The  following  are 
interesting  examples : 

Charcoal  glows  in  air  only  moderately  and  without  much  evolu^ 
tion  of  light.  In  oxygen,  however,  it  burns  with  a  bright  glow. 
Sulphur,  which  burns  in  air  with  only  a  small  flame,  burns  in 
oxygen  with  an  intense  blue  light.  Phosphorus  burns  in  oxygen 


12  INORGANIC  CHEMISTRY.  [§§  10- 

with  a  blinding  white  light.  A  steel  watch-spring  that  has  been 
heated  to  redness  at  one  end  and  put  into  oxygen,  burns  with 
scintillation.  Zinc  also  burns  in  it  with  a  dazzling  light.  In  all 
these  and  analogous  cases  the  oxygen,  as  well  as  the  burning  mate- 
rial, disappears  during  the  combustion,  while  new  substances  are 
formed.  A  lasting  change  therefore  takes  place  and  we  have  to 
do  with  a  chemical  process.  The  product  of  burning  charcoal  is 
found  to  be  a  gas  that  makes  lime-water  cloudy  and  is  unable  to 
support  combustion;  it  is  called  carbonic  acid  gas.  Sulphur  also 
yields  a  gas;  it  has  a  pungent  odor  and  is  called  sulphur  dioxide. 
Phosphorus  produces  a  white  flocculent  powder,  phosphorus  pent- 
oxide.  When  iron  burns,  a  black  cindery  powder  is  formed, 
called  "hammer-scale,"  because  it  composes  the  sparks  that  fly 
from  the  anvil. 

The  question  now  arises  as  to  what  really  occurs  in  the  above 
cases.  In  the  first  place,  it  has  been  found  that  the  weight  of  the 
product  of  combustion  is  greater  than  that  of  the  substance  burned. 

The  increase  in  weight  of  the  substance  during  burning  can  in  many 
cases  be  easily  demonstrated.  For  instance,  a  horseshoe  magnet  that  has 
been  dipped  in  iron  filings  may  be  hung  on  the  lower  side  of  a  scale-pan 
and  balanced  by  weights  put  in  the  other  pan.  The  iron  filings  may  be 
burned  by  passing  a  non-luminous  flame  under  them  a  few  times.  On 
cooling,  the  scale-pan  attached  to  the  magnet  sinks.  In  a  similar  way  one 
may  demonstrate  the  increase  of  weight  in  the  burning  of  copper.  In 
order  to  prove  the  increase  of  weight  in  a  case  where  only  gaseous  products 
are  formed,  a  candle  may  be  burned  and  the  combustion  products, 
carbon  dioxide  and  water  vapor,  collected  by  letting  them  pass  over 
unslaked  lime,  with  which  both  unite. 

Closer  investigation  has  revealed  the  fact  that  the  increased 
weight  is  due  to  the  presence  of  oxygen,  as  well  as  the  burned 
substance,  in  all  combustion  products.  The  latter  are  compounds 
of  these  substances  with  oxygen.  The  participation  of  oxygen  in 
the  burning  of  zinc,  for  example,  may  be  proved  by  heating  the 
combustion  product,  zinc  white,  in  a  tube  and  leading  over  it 
chlorine  gas,  whereby  oxygen  is  driven  oft7.  The  compounds  of 
oxygen  are  called  oxides,  and  the  act  of  this  combination  is  known 
as  oxidation. 

When  substances  burn  in  the  air,  it  is  only  the  oxygen  which 
combines  with  them.  Nevertheless,  the  nitrogen  of  the  air  is  heated 


11.]  HYDROGEN.  13 

and  thus  takes  a  part  of  the  heat  evolved  in  the  combustion. 
Therefore  the  temperature  of  a  burning  object  cannot  rise  so  high 
as  in  pure  oxygen,  and,  since  the  emission  of  light  increases  very 
rapidly  as  the  temperature  rises,  combustions  in  oxygen  are  for 
this  reason  much  brighter  than  in  air. 

There  are  two  general  methods  of  ascertaining  what  elements 
exist  in  a  compound.  According  to  the  one  method  the  compound 
is  decomposed  and  the  elements  composing  it  thereupon  deter- 
mined. This  is  the  analytic  method.  According  to  the  other, 
the  synthetic  method,  the  composition  is  found  by  combining 
different  elements  to  form  new  substances.  In  the  above-described 
experiment  (§  8)  of  LAVOISIER  the  composition  of  the  red  powder 
is  learned  by  decomposing  it  at  high  temperature,  whereupon  it 
separates  into  only  mercury  and  oxygen.  Inversely  it  was  possible 
to  obtain  the  red  powder  by  heating  pure  oxygen  and  pure  mercury 
together  at  a  lower  temperature.  The  former  is  an  example  of 
analysis,  the  latter  of  synthesis. 

HYDROGEN. 

ii.  Hydrogen  is  a  colorless  and  odorless  gas  that  is  rarely 
found  on  the  earth  in  the  free  state.  The  gases  of  some  volcanoes 
contain  it  and  it  can  also  result  from  processes  of  decay.  In 
combination  with  other  elements,  however,  hydrogen  is  very 
widely  distributed  and  occurs  in  very  large  amounts  (§8). 

Hydrogen  can  be  prepared  in  various  ways.  In  the  first  place, 
hydrogen  compounds  can  be  broken  up. 

(1)  Water  containing  some  dissolved  electrolyte  is  decomposed 
by  the  electric  current  with  the  evolution  of  hydrogen  at   the 
negative  pole  (cathode). 

The  ordinary  methods  of  preparing  hydrogen  depend  on  the 
indirect  decomposition  of  hydrogen  compounds,  i.e.,  their  reaction 
with  other  substances.  The  following  are  examples  of  this  sort: 

(2)  The  action  of  zinc  on  dilute  sulphuric  acid  (§  89).     This 
is  the  commonest  method. 

For  the  preparation  of  hydrogen  in  the  laboratory  the  apparatus 
shown  in  Fig.  5  is  often  used.  A  contains  granulated  zinc  (or  iron  nails) 
and  B  dilute  hydrochloric  acid  or  sulphuric  acid.  When  the  cock  C  is 
opened  the  acid  flows  through  D  to  the  metal  and  the  evolution  of  hy- 
drogen commences  at  once.  The  cock  being  closed  again,  the  gas  still 


14 


INORGANIC  CHEMISTRY. 


[§§  11- 


continues  to  come  off  and  forces  back  the  acid.  This  is  facilitated  by- 
changing  the  relative  levels  of  A  and  B. 

(3)  The  action  of 
zinc  or  aluminium  fil- 
ings on  caustic  potash 
or  slaked  lime. 

(4)  The  action  of 
sodium  or  potassium 
on  water  or  alcohol. 

(5)  Magnesium 
powder,  when  boiled 
with      water,      also 
evolves       hydrogen, 
especially  when  some 

chloride    of    magne-  FlG* 5t 

sium  is  dissolved  in  the  water,  because  such  a  solution  dissolves 
the  magnesium  oxide  which  forms  on  the  surface  of  the  metal. 
Likewise,  red-hot  iron  decomposes  water  with  the  liberation  of 
hydrogen  (compare  §  305). 

i  12.  The  physical  properties  of  hydrogen  are  these:  It  is  the 
lightest  of  all  known  substances,  its  specific  gravity  (air=  1) 
amounting  to  only  0.06949.  One  liter  of  hydrogen  at  0°  and 
760  mm.  Hg.  pressure  weighs  0.0899  g.  Its  lightness  renders  it 
useful  for  inflating  balloons.  It  is  very  hard  to  liquefy,  because 
its  critical  temperature  lies  only  30-32°  above  the  absolute  zero 
(—273°).  On  the  other  hand,  the  critical  pressure  is  only  15 
atmospheres.  Liquid  hydrogen  is  colorless.  It  boils  at  —  252.5°. 
Its  specific  gravity,  with  reference  to  water,  is  only  0.07  at  its 
boiling-point  and  0.086  at  its  freezing-point,  being  therefore 
considerably  less  than  that  of  all  other  known  liquids.  DEWAB 
further  succeeded  in  bringing  hydrogen  to  the  solid  state  by 
allowing  the  liquid  to  evaporate  quickly  at  30-40  mm.  pressure. 
The  melting-point  of  solid  hydrogen  is  about  16°  (absolute  tem- 
perature). The  heat  of  evaporation  of  liquid  hydrogen  is  very 
high,  being  200  cal. ;  for  this  reason  a  flask  containing  liquid 
hydrogen  soon  becomes  Covered  with  a  layer  of  liquid  air,  which 
drops  down  and  soon  partially  solidifies. 

Hydrogen  is  slightly  soluble  in  water,  100  1.  water  dissolving 
2.15  1.  of  the  gas  at  0°.     Alcohol  takes  up  somewhat  more. 

13.  Chemical  Properties. — Hydrogen  does  not  unite  with  as 
large  a  number  of  elements  as  oxygen.     At  a  higher  temperature 


13.]  HYDROGEN.  15 

it  displays  a  strong  tendency  to  unite  with  oxygen,  burning  with 
an  almost  colorless  and  a  very  hot  flame  to  form  water.  This 
property  serves  for  the  identification  of  hydrogen  gas. 

When  a  current  of  hydrogen  is  directed  upon  very  finely 
divided  platinum  (spongy  platinum  or  platinum  black,  §  316), 
the  hydrogen  is  ignited  (§  25). 

The  high  temperature  of  the  hydrogen  flame  is  made  use  of  in 
fusing  platinum,  quartz,  etc. 

Such  a  flame  is  known  as  an  oxyhydrogen  flame.  An  apparatus 
(oxyhydrogen  bloivpipe)  like  that  represented  in  Fig.  6  is  required  for 
producing  it.  The  hydrogen  enters  at  W  and  passes  out  at  a,  where  it  is 
lit.  Oxygen  is  blown  into  the  flame  at  S.  Thus  the  gases  do  not  mix 
till  they  reach  the  flame,  and  the  possibility  of  an  explosion  is  avoided. 


FIG.  0. — OXYHYDROGEN  BLOWPIPE. 

A  mixture  of  hydrogen  and  oxygen,  especially  in  the  proportion 
of  2  vols.  H  and  1  vol.  O  (detonating-gas),  when  ignited,  turns 
instantaneously  to  steam;  in  other  words,  it  explodes.  This  ex- 
periment can,  however,  be  performed  harmlessly  by  using  a  wide- 
mouthed  cylinder  of  not  too  great  dimensions.  A  loud  report  is 
heard  in  this  case,  because  the  steam  at  the  moment  of  its  forma- 
tion occupies  a  much  larger  volume  at  the  high  temperature  of  the 
combustion  than  the  mixture  of  the  original  gases,  and  as  a  result 
the  air  is  suddenly  ejected  with  violence.  When  the  explosion 
occurs  in  a  closed  vessel,  no  sound  is  heard  (cf.  e.g.  Fig.  13,  p.  25). 

The  temperature  to  which  detonating-gas  must  be  heated  to 
explode  is  found  to  be  about  700°.  At  a  lower  temperature  com- 
bination between  hydrogen  and  oxgyen  also  takes  place,  but  not 
instantaneously,  as  in  explosions;  the  lower  the  temperature,  the 
slower  the  process.  When,  therefore,  no  change  in  cold  detonating- 
gas  is  observed  even  in  the  course  of  several  years,  we  must  attribute 
the  fact  to  the  extraordinary  slowness  of  the  process  at  ordinary 
temperatures.  A  simple  calculation  will  make  this  plain. 
BODENSTEIN  observed  that,  when  detonating-gas  is  heated  at 
509°  for  50  minutes,  0.15  of  the  whole  is  changed  to  water.  Now 
it  is  a  general  rule  that,  when  the  temperature  sinks  10°,  a  chemi- 
cal reaction  becomes  about  twice  as  slow;  at  499°  it  would  thus 


16  INORGANIC  CHEMISTRY.  [§§  13- 

take  100  minutes  till  the  0.15  part  of  the  gas  had  formed  water. 
At  the  ordinary  temperature,  say  at  9°,  it  would  be  50 X250  minutes, 
that  is  about  1.06X  1011  years.  The  same  can  be  said  of  all  chemi- 
cal reactions.  When  we  see  that  wood,  sulphur,  etc.,  burn  quickly 
at  higher  temperatures,  we  must  admit  that  oxidation  takes 
place  also  at  ordinary  temperatures,  though  so  slowly  that  we 
cannot  perceive  it.  MOISSAN,  however,  succeeded  in  proving 
that  charcoal  at  100°  and  sulphur  at  ordinary  temperatures  are 
oxidized  very  slowly  in  a  current  of  oxygen. 

Hydrogen  is  not  only  able  to  unite  with  free  oxygen,  but  it  also 
has  the  power  to  withdraw  oxygen  from  many  of  its  compounds. 
The  action  of  hydrogen  on  a  compound  is  called,  in  general, 
reduction.  This  action  is  often  a  very  useful  means  of  determining 
whether  a  compound  contains  oxygen,  since  the  latter,  if  present, 
frill  usually  unite  with  the  hydrogen  to  form  water.  Copper  oxide 
may  serve  as  an  example  of  the  application  of  this  method.  A 
little  is  placed  in  a  tube,  hydrogen  is  led  over  it,  and  heat  is  then 
applied ;  one  soon  sees  the  black  oxide  change  to  red  copper,  and 
water  depositing  in  drops  on  the  colder  parts  of  the  tube.  Many 
other  oxides  can  be  similarly  reduced,  e.g.  iron  oxide,  lead  oxide, 
etc. 

THE  CONSERVATION   OF  MATTER. 

14.  The  quantitative  relationships  in  oxidizing  and  reducing 
processes,  such  as  have  been  discussed  in  §  13,  i.e.  the  relations 
of  the  masses  of  the  substances  participating  in  the  changes,  may 
be  used  to  elucidate  a  very  important  law.  A  definite  amount  of 
copper  powder,  for  example,  may  be  placed  in  a  tube  and  the 
weight  of  the  tube  with  the  powder  ascertained.  Oxygen  is  then 
led  over  the  copper  at  a  high  temperature.  The  apparatus  should 
be  so  arranged  that  the  volume  of  the  oxygen  which  combines  with 
the  copper  can  be  measured.  When  the  oxidation  process  has 
proceeded  for  some  time,  the  tube  containing  the  oxidized  copper 
is  allowed  to  cool  and  then  weighed.  The  weight  is  found  to  have 
increased,  and  the  increase  is  just  equal  to  the  weight  of  the  volume 
of  oxygen  used  up.  Thereupon  hydrogen  is  passed  through  the 
tube  with  the  copper  oxide  and  heat  applied.  Here  also  arrange- 
ments should  be  made  for  measuring  the  volume  of  hydrogen  con- 
sumed in  reduction.  The  reduction  is  allowed  to  go  on  until  all  the 
copper  oxide  is  transformed  back  to  copper.  When  the  tube  and 
powder  are  subsequently  weighed,  they  will  be  found  to  have  re- 


15.]  WATER.  17 

assumed  their  original  weight.  The  water  that  forms  can  be 
absorbed  by  a  substance  like  quicklime  or  concentrated  sulphuric 
acid  and  weighed.  It  will  be  found  equal  in  weight  to  the  loss  of 
weight  of  the  copper  oxide  on  changing  to  copper  plus  the  weight 
of  the  consumed  hydrogen. 

In  these  cases,  therefore,  the  combined  weight  of  the  reacting 
substances  before  and  after  the  reaction  is  the  same.  Copper + 
consumed  oxygen  weighs  just  as  much  as  copper  oxide;  copper 
oxide -f  consumed  hydrogen  weighs  just  as  much  as  copper + water; 
and,  finally,  the  regained  copper  weighs  just  as  much  as  that  origi- 
nally taken.  The  substances  can  be  changed  into  different  states, 
but  their  weight  remains  unaltered.  This  phenomenon  is  observed 
without  exception  in  chemical  actions,  and  we  therefore  accept  as  a 
law  the  statement  that  matter  is  indestructible,  or  that  no  matter 
can  be  lost  or  gained.  This  principle  was  introduced  into  chemistry 
by  LAVOISIER  (1743-1794;. 

The  old  Greek  philosophers  were  already  firmly  convinced  of  the 
impossibility  of  producing  or  destroying  matter.  In  all  ages  this  belief  has 
been  the  basis  of  philosophic  thought.  To  LAVOISIER  is  due  the  credit 
of  having  demonstrated  the  practical  application  of  the  principle  of  the 
indestructibility  of  matter.  He  assumed  that  gravity  is  an  inseparable 
attribute  of  all  matter — concerning  which  a  great  deal  of  doubt  still 
existed — and  that  the  combined  weight  of  the  substances  concerned 
must  therefore  be  the  same  before  and  after  a  chemical  reaction. 

The  theory  of  knowledge  teaches  that  the  principle  of  the  indestructi- 
bility of  matter  lies  originally  at  the  basis  of  our  thinking.  It  is  entirely 
incorrect  to  suppose  that  it  was  established  by  experimentation;  on  the 
contrary,  we  test  the  correctness  of  our  experimental  results  by  ascertain- 
ing in  how  far  they  conform  to  this  principle.  This  can  be  easily  under- 
stood in  the  above  case  of  the  oxidation  and  reduction  of  copper.  In  per- 
forming this  experiment  one  finds  that  the  weight  of  copper  +  oxygen  is 
not  exactly  equal  to  that  of  the  copper  oxide  formed.  Even  after  several 
repetitions  slight  differences  are  still  found.  Because  we  feel  that  there 
must  be  absolute  equality,  we  attribute  these  differences  to  imperfections 
in  our  instruments,  and  we  consider  our  instruments  improved  if  they 
enable  us  to  approach  nearer  the  complete  equality  of  the  weights  before 
and  after  the  experiment.  Nevertheless,  we  are  unable  to  really  observe 
an  absolute  equality. 

WATER. 

15.  Water  was  regarded  as  an  element  for  many  centuries.  Not 
until  1781  did  CAVENDISH  discover  that,  when  a  mixture  of  hydro- 


18 


INORGANIC  CHEMISTRY. 


[§15. 


gen  and  air  or  oxygen  explodes,  water  is  formed.  Being,  how- 
ever, a  supporter  of  an  erroneous  theory  (§  106),  he  failed  to  realize 
the  importance  of  his  discovery.  LAVOISIER  in  1783  repeated  this 
experiment  and  comprehended  it  as  a  synthesis  of  water,  as  we  still 
do  to-day. 

With  the  aid  of  the  apparatus  pictured  in  Fig.  7,  this  synthesis 
can  be  easily  demonstrated.    The  hydrogen  is  generated  in  the 


FIG.  7. — COMBUSTION  OF  HYDROGEN. 

two-necked  (WOULFP)  bottle  from  zinc  and  sulphuric  acid.  In 
order  to  free  the  gas  from  water  vapor,  it  is  passed  through  the 
horizontal  tube,  which  contains  chloride  of  calcium  or  bits  of 
pumice-stone  soaked  in  sulphuric  acid.  The  dry  gas  is  ignited  and, 
as  it  burns,  water  is  gradually  deposited  on  the  walls  of  the  bell -jar. 
A  mixture  of  hydrogen  and  oxygen  unites  to  form  water  when  illu- 
minated with  ultraviolet  light. 

In  addition  to  this  direct  synthesis  from  its  elements  there  are 
other  ways  of  obtaining  water.  For  example,  many  compounds, 
such  as  the  blue  crystals  of  copper  vitriol,  give  off  water  when 
heated. 

The  formation  of  water  by  the  action  of  hydrogen  on  oxygen 
compounds  was  illustrated  (§  13)  in  the  reduction  of  copper  oxide. 
On  the  other  hand,  it  is  also  produced  by  the  action  of  oxygen  on 
certain  hydrogen  compounds.  This  is  seen,  for  example,  in  the 
burning  of  alcohol. 

Finally,  water  can  result  from  the  reaction  of  a  hydrogen  com- 
pound with  one  of  oxygen.  This  is  the  case  when  ammonia  gas 
(§  111)  is  led  over  hot  copper  oxide. 


16. 


WATER. 


19 


8 


The  synthetic  methods  of  preparing  water,  such  as  the  above- 
named  and  many  others,  possess,  however,  merely  theoretical 
importance.  Even  when  water  is  wanted  in  a  perfectly  pure 
state,  natural  water  is  resorted  to.  This  contains  solids  and  gases 
in  solution,  which  must  be  eliminated.  Its  purification  is  accom- 
plished by  distillation.  An  apparatus  well  suited  to  this  purpose 
is  shown  in  Fig.  8.  High  pressure  steam  and  electricity  are 
often  used  for  heating  instead  of  the  flame. 

Water  is  placed 
in  the  retort  A, 
which  rests  over  the 
fireplace,  and  boil- 
ed. The  dissolved 
gases  are  first  driven 
off;  the  hot  steam 
follows,  passing 
t  hrough  the  dome  B 
into  the  condensing 
coil  ("  worm  ")  C, 
which  is  cooled  by 
water  in  the  vessel 
Z).  The  condensed 
water,  now  pure, 
flows  down  into  the 

bottle;  the  solid PIG>  s.— PURIFICATION  OF  WATER  BY  DISTILLATION. 
substances  that 

were  dissolved  in  the  water  remain  in  the  retort.  The  cooler 
D  is  supplied  with  cold  water  through  a  tube,  entering  near 
the  bottom,  while  the  heated,  and  therefore  specifically  lighter, 
water  flows  out  near  the  top.  The  steam  thus  meets  with  cooling- 
water  of  a  lower  temperature  as  it  passes  down  the  worm,and  is  in  this 
way  very  completely  condensed  (principle  of  the  counter-current). 

A  single  distillation  is  usually  insufficient  for  the  complete  elimination 
of  all  gaseous  and  solid  constituents.  For  this  purpose  the  operation 
must  be  repeated  in  an  apparatus  of  platinum  (tin  is  less  satisfactory) 
with  a  condensing  coil  of  the  same  metal,  and  only  the  middle  fraction 
collected. 

An  excellent  criterion  for  the  purity  of  water  is  to  be  found  in  the 
measurement  of  its  electrical  resistance.  Very  pure  water  conducts  the 
electric  current  scarcely  at  all.  KOHLRAUSCH  found  the  conductivity  at 


20  INORGANIC  CHEMISTRY.  [§§15- 

18°  of  the  purest  water  obtainable  to  be  k  =  0.038 X10~6  expressed  in 
reciprocal  ohms;  by  this  is  meant  the  conductivity  of  a  body  a  column 
of  which  1  cm.  long  and  1  cm.  square  in  cross-section  has  a  resistance 
of  1  ohm.  The  magnitude  of  the  resistance  of  such  water  is  better  un- 
derstood by  comparing  it  with  resistance  of  copper.  1  cu.  mm.  of  this 
water  has  at  0°  the  same  resistance  as  a  copper  wire  of  the  same  cross- 
section  and  25  million  miles  long;  it  could  be  strung  around  the  earth's 
equator  one  thousand  times.  The  slightest  traces  of  salts  or  even  con- 
tact with  the  atmosphere  cause  a  market  increase  in  its  conductivity. 

PHYSICAL  PROPERTIES. 

1 6.  Water  at  ordinary  temperatures  is  an  odorless,  tasteless 
liquid,  showing  no  color  in  thin  layers.     On  looking  through  a  layer 
26  meters  thick,  SPRING  observed  a  pure  dark-blue  color.     The 
thermometer-scale  of  CELSIUS  is  fixed  according  to  the  physical 
constants  of  water,  its  freezing-point  being  called  0°  and  its  boiling- 
point  at  760  mm.  pressure  100°.     These  two  points  are  dependent 
on  the  pressure.     An  increase  of  pressure  lowers  the  freezing-point 
(0.0075°  per  atmosphere).     This  is  the  reason  why  ice  melts  under 
high  pressure.     Water  possesses  the  very  uncommon  property  of 
having  a  maximum  of  density  (minimum  of  volume)  at  a  definite 
temperature.    The  volume  of  almost  all  other  substances  increases 
with  rising  temperature,  but  here  it  diminishes  up  to  3.945°,  above 
which  temperature  water  expands  as  heating  continues.     During 
the  transformation  of  water  to  ice  the  volume  increases  considerably. 
One  vol.  water  at  0°  yields  1.09082  vol.  ice  of  the  same  temperature. 

The  specific  heat  of  water  is  greater  than  that  of  a  vast  majority 
of  other  substances.  Its  latent  heat  of  fusion  is  79  Cal.,  its  latent 
heat  of  vaporization  536  Cal.  Water  is  extensively  used  as  a  sol- 
vent. Numerous  substances  dissolve  in  it  to  a  greater  or  less  degree. 
There  are  many  liquid  substances  that  mix  with  water  in  all  pro- 
portions, and  many,  also,  which  do  not.  (See  §  7.) 

The  remarkable  physical  properties  of  water  play  a  very  important 
role  in  nature;  this  subject  is  extensively  discussed  in  physics,  meteor- 
ology, and  geology. 

NATURAL  WATER. 

17.  Water,  as  it  occurs  in  nature,  is  by  no  means  chemically 
pure.     It  may  contain  solid  matter  in  suspension  as  well  as  sub- 
stances, either  solid  or  gaseous,  in  solution.     The  purest  natural 
water  is  rain-water.     This  has  really  passed  through  a  natural 
process  of  distillation,  the  water  on  the  earth's  surface  being  vapor- 


17.] 


NATURAL  WATER. 


21 


ized  by  the  sun's  heat  and  condensed  again  by  contact  with  colder 
portions  of  air,  whereupon  it  falls  in  the  form  of  rain.  Neverthe- 
less it  contains  dust  particles  (in  large  cities  more,  of  course,  than 
in  the  country)  and  gases  from  the  air,  as  well  as  traces  of  ammo- 
nium salts. 

Spring-  and  well-waters  contain  in  10,000  parts  about  1-20  parts 
of  solid  matter,  consisting  largely  of  lime  salts.  Well-water  that  con- 
tains much  lime  is  called  hard  (§  259).  Well-water  also  contains 
some  carbonic  acid  and  air  in  solution,  both  of  which  give  it  its 
refreshing  taste;  distilled  water  tastes  flat. 

Natural  water  is  used  extensively  for  drinking  purposes.  When  it 
comes  out  of  a  soil  that  is  contaminated  by  decaying  organic  matter,  as  is 
the  case  in  many  large  cities,  it  is  injurious  to  health,  principally  on 
account  of  the  presence  of  bacteria.  It  can  be  freed  from  these  by 
filtration  through  a  PASTEUR-CHAMBERLAND  porcelain  filter  (Fig.  9). 

This  consists  essentially  of  a  hollow  cylinder  of  porous  porcelain  (called 
a  "candle")  A,  through  whose  walls  the  water  is  forced  by  its  own  pres- 
sure. The  lower  end  of  the  candle  opens  into 
the  nozzle. 

In  large  cities  it  has  been  found  much  more 
practicable  to  purify  the  well-  or  river-water  at 
the  central  station  and  to  pipe  it  thence  to  the 
various  houses.  Epidemic  diseases  have  really 
decreased  remarkably  since  the  introduction  of 
the  methods  of  modern  sanitary  science. 

A  water  which  contains  so  many  substances  in 
solution  that  it  has  a  definite  taste  or  a  therapeu- 
tic effect  is  called  a  mineral  water.  There  are  very 
many  kinds  of  mineral  waters,  differing  accord- 
ing to  the  amount  and  kind  of  dissolved  matter 
they  contain.  We  distinguish  between  saline 
waters  containing  common  salt,  bitter 
waters  with  magnesium  salts,  sulphu- 
rous waters  with  sulphuretted  hydrogen, 
carbonated  waters  with  carbonic 
acid,  chalybeate  waters  with  iron, 
and  many  others.  Detailed  analyses  of  the 
mineral  waters  of  numerous  watering-places  are  accessible  in 
works  on  balneology. 

Sea-water  contains  about  3%  of  salts,  of  which  2.7%  is  common 


FIG.  9. — PASTEUR- 
CHAMBERLAND 
FILTER. 


22 


INORGANIC  CHEMISTRY. 


[§§  17- 


salt.  A  large  number  of  elements,  viz.,  about  thirty,  have  been 
found  in  sea-water,  although  the  most  of  them  exist  there  only  in 
extremely  small  quantities. 

It  was  stated  above  (§16)  that  pure  water  is  blue.  The  color  of  the 
rivers,  lakes  and  seas  varies,  however,  through  many  nuances  from  pure 
blue  to  brown.  This  variation  is  due  principally  to  the  presence  of  more 
or  less  brownish-yellow  humous  (marshy)  substances  or  an  extremely 
fine  floating  slime.  Both  conditions  can  produce  a  brownish-yellow 
color.  It  is  easily  seen  how  the  combination  of  blue  and  yellow  or  brown 
may  bring  about  the  various  blue,  green  or  brown  tints  in  natural  waters. 


COMPOSITION  OF  WATER. 

18.  Decomposition. — It  was  stated  above  that  water  can  be 
obtained  by  direct  combination  of  hydrogen  and  oxygen;  inversely, 
it  can  be  decomposed  into  these  same  elements. 

In  the  flask  A  (Fig.  10)  some  water  is  heated  till  it  boils  vigorously.  A 
strong  electric  current  is  then  sent  through  the  wire  a  c  b,  so  that  the  fine 
platinum  wire  c  glows  intensely.  This  heat  partially  decomposes  the 


f                          ^ 

?        'f        'II         ^ 

^                         !M                              If, 

II"                    -'||. 

FIG.  10. — DECOMPOSITION  OF  WATER  BY  GLOWING  PLATINUM, 

water  vapor  into  hydrogen  and  oxygen,  which  pass  out  through  the  tube 
d  and  are  collected  in  the  cylinder  C.  This  gas  mixture  is  nothing  but  the 
explosive  mixture  (§  13)  of  hydrogen  and  oxygen,  as  can  be  easily  proved 
by  applying  a  flame. 


19.] 


COMPOSITION  OF  WATER. 


23 


Many  metals  decompose  water  on  contact,  the  hydrogen  being 
set  free  and  the  metal  uniting  with  the  oxygen.  Potassium  and 
sodium  effect  this  decomposition  at  ordinary  temperatures  (§  11); 
iron,  zinc  and  other  metals  require  a  higher  temperature,  iron, 
e.g.,  acting  at  red  heat. 

19.  Let  us  now  study  the  quantitative  composition  of  water, 
i.e.  determine  the  relative  amounts  of  hydrogen  and  oxygen  present. 
For  this  purpose  both  the  analytic  and  synthetic  methods  can  be 
used. 

(a)  The  Analytic  Method. — When  an  electric  current  is  passed 
through  water  to  which  has  been  added  a  little  sulphuric  acid,  the 
water  is  decomposed.  If  the  gases  evolved  at  the  electrodes  are 
collected  separately,  it  is  found  that  for  every  1  vol.  oxygen  2  vols. 
hydrogen  are  given  off.  A  suitable  apparatus  for  this  experiment 
is  shown  in  Fig.  11. 

Since  1  liter  of  hydrogen  weighs  0.0899  g.  and  1  liter  of  oxygen 
weighs  1.4296  g.,  both  at  0°  and  760  mm. 
pressure,  the  weights  of  2  vols.  hydrogen 
and  1  vol.  oxygen  must  bear  to  each 
other  the  ratio  of  2X0.0899  :  1.4296,  or 
1  :  7.943. 

(6)  The  Synthetic  Method.— As  early 
as  1820  the  reduction  of  copper  oxide  by 
hydrogen  was  employed  for  this  purpose 
by  BERZELIUS;  in  1834,  also,  by  DUMAS 
and  STAS.  A  weighed  amount  of  care- 
fully dried  copper  oxide  is  heated  in  a 
current  of  hydrogen  and  water  is  formed, 
which  is  collected  and  weighed.  The 
weight  of  the  oxygen  given  up  by  the 
copper  oxide  is  found  from  the  difference 
between  the  weight  of  the  copper  oxide 
used  and  that  of  the  resulting  copper. 
The  weight  of  the  hydrogen  contained 
in  the  water  collected  is  therefore  equal 
to  the  difference  in  weight  of  water  and 
oxygen. 

The  apparatus  used  for  this  experiment  is  represented  in  Fig.  12. 
In  A  the  hydrogen  is  generated  from  zinc  and  dilute  sulphuric 


FIG.  11. — ELECTROLYSIS 
OF  WATER. 


24 


INORGANIC  CHEMISTRY. 


19- 


acid.  It  is  then  passed  through  the  permanganate  solution  in  the 
wash-bottle  B  to  free  it  from  impurities,  and  also  through  the  U- 
tubes  C,  D  and  E,  containing  calcium  chloride,  sulphuric  acid  and 
phosphorus  pentoxide,  respectively,  for  drying  it.  In  F  is  placed 
the  copper  oxide,  which  is  carefully  weighed  together  with  the  tube. 
The  water  that  forms  is  condensed  in  G,  the  U-tube  H  being 
attached  to  absorb  any  escaping  water  vapor.  At  the  completion 
of  the  experiment,  F,  with  its  contents,  is  again  weighed,  likewise 
G  and  H;  the  differences  in  weight  indicate  the  amount  of  water 


FIG.  12. — SYNTHESIS  OF  WATER  AFTER  DUMAS  AND  STAS. 

formed.  DUMAS  and  STAS  found  in  this  way  that  100  parts  (by 
weight)  of  water  consist  of  11.136  parts  of  hydrogen  and  88.864 
parts  of  oxygen,  or,  in  other  words,  that  the  mass-ratio  of  these 
elements  is  1:7.980,  a  relation  which  agrees  with  that  ob- 
tained in  (a)  within  the  range  of  the  unavoidable  experimental 
error. 

Another  synthetic  method,  which  is  especially  adapted  to  the 
lecture-table,  consists  in  mixing  hydrogen  and  oxygen  and  deter- 
mining in  what  volume-ratio  these  gases  unite.  For  this  pur- 
pose an  apparatus  (Fig.  13)  described  by  HOFMANN  is  best  em- 
ployed. 

Hydrogen  and  oxygen  in  different  proportions  by  volume  are 
introduced  into  the  arm  of  the  U-tube,  which  can  be  closed  by 


20.] 


COMPOUNDS  AND  MIXTURES. 


25 


a  stop-cock  at  the  top :  the  cock  is  thereupon  closed  and  the 
arm  tightly  stoppered  with  a  cork.  The 
mixture  is  then  exploded  by  an  in- 
duction spark,  the  volume  of  air  en- 
closed on  the  other  side  acting  as  a 
cushion  to  moderate  the  severe  shock 
on  the  mercury,  which  might  otherwise 
break  the  apparatus.  It  is  found  that 
only  when  the  volumes  of  hydrogen 
and  oxygen  bear  to  each  other  the  ratio 
2:1  does  the  entire  gas  mixture  dis- 
appear, a  slight  coating  of  tiny  drops  of 
water  appearing  in  its  place  on  the 
inside  of  the  glass.  In  case  more  hy- 
drogen or  more  oxygen  than  the  ratio 
calls  for  is  let  into  the  tube,  the  excess 
is  found  to  remain  after  the  explo- 
sion. 

From  these  experiments,  analytic  and 
synthetic,    it   follows   that  water  has  a  FIG.   13. — HOFMANN'S 
constant  composition;  it  consists  of  2  vols. 
of  hydrogen  and   1  vol.  of  oxygen,  or 
of  1  part,  by  weight,  of  hydrogen  to  7.943  parts  of  oxygen. 


open 


AP- 

SYN- 


COMPOUNDS  AND  MIXTURES. 

20.  In  water  we  have  become  acquainted  with  a  substance  which 
is  different  in  many  and  important  respects  from  the  elements  of 
which  it  is  composed.  We  have  further  seen  that  the  elements  in 
it  bear  to  each  other  a  fixed  relation  by  weight.  Such  substances 
are  known  in  very  large  number.  Copper  oxide,  mercury  oxide, 
sulphuric  acid,  potassium  chlorate,  common  salt,  soda  and  many 
others  already  mentioned  belong  to  this  class.  In  each  of  these, 
no  matter  how  obtained,  we  discover  by  analysis  or  synthesis  a 
definite  proportion  between  the  elements  composing  it.  Such  sub- 
stances are  called  compounds. 

In  addition  to  the  characteristics  mentioned — difference  of  prop- 
erties from  those  of  the  elements  and  constant  composition — we 
find  that  the  compounds  also  have  constant  physical  properties. 


26  INORGANIC  CHEMISTRY.  [§§  20- 

Under  the  same  pressure  water  always  has  the  same  melting-point 
and  the  same  boiling-point,  in  whatsoever  way  it  may  have  been 
obtained;  salt  always  crystallizes  in  the  same  crystal  system;  soda, 
at  a  definite  temperature,  always  requires  the  same  amount  of  water 
for  solution,  etc. 

When  elements  or  compounds  are  brought  together  without  any 
chemical  action  on  each  other  taking  place,  we  have  a  mixture  of 
these  elements  or  compounds.  The  number  of  possible  mixtures  is, 
of  course,  unlimited.  They  are  distinguished  from  compounds  by 
the  following  characteristics : 

In  a  mixture  the  properties  of  the  components  reappear  in  many 
and  important  respects.  Gunpowder,  for  example,  is  a  mixture  of 
sulphur,  charcoal  and  saltpetre.  The  latter  is  soluble  in  water; 
sulphur  dissolves  in  carbon  disulphide;  charcoal  is  insoluble  in  both. 
These  properties  are  still  evident  in  the  constituents  of  gunpowder. 
In  a  mixture  of  sulphur  and  iron  filings  one  can  detect  with  a  micro- 
scope the  yellow  grains  of  sulphur  and  the  black  particles  of  iron. 
The  iron  can  be  drawn  out  with  a  magnet;  the  sulphur  dissolved 
out  by  carbon  disulphide.  If,  however,  a  mixture  of  7  parts  iron 
and  4  parts  sulphur  is  heated,  a  glow  passes  through  the  powder 
and  a  compound  of  both — iron  sulphide — is  formed,  whose  prop- 
erties are  entirely  different  from  those  of  its  elements.  It  is  non- 
magnetic and  insoluble  in  carbon  disulphide  and  under  the  micro- 
scope only  a  homogeneous  scoriaceous  mass  is  seen.  The  constituents 
of  a  mixture,  since  they  still  preserve  their  properties,  can  often  be 
separated  from  each  other  by  mechanical  means,  e.g.  by  the  use  of 
microscope  and  tweezers,  by  sifting,  by  treatment  with  solvents,  by 
washing,  etc. 

In  a  mixture  the  ratio  of  the  constituents  can  vary  in  all  pro- 
portions. There  are,  for  example,  many  sorts  of  gunpowder,  dis- 
tinguished from  each  other  by  the  proportions  in  which  their  con- 
stituents are  mixed.  When  1  part  sulphur  and  100  parts  iron,  or, 
on  the  other  hand,  1  part  iron  and  100  parts  sulphur,  are  mixed,  we 
have  in  either  case  a  mixture  of  both  elements,  possessing  hardly 
the  same,  but  at  least  analogous,  properties. 

Moreover,  a  mixture  often  has  no  constant  physical  properties. 
Water  has  a  constant  boiling-point;  the  boiling-point  of  a  mixture 
of  benzene  and  turpentine,  however,  rises  gradually  as  the  more 
volatile  component,  benzene,  distils  off.  The  melting-point  of 


21,]       COMPOSITION  OF  COMPOUNDS.— ATOMIC  THEORY.       27 

sulphur  is  constant  and  can  be  accurately  determined;  that  of  a 
mixture  of  tin  and  lead  differs  according  to  the  proportion  of  the 
elements  and  is  in  many  proportions  not  at  all  sharp,  there  being 
only  a  softening  instead  of  real  fusion. 

In  the  examples  cited  here  the  distinction  between  a  compound  and  a 
mixture  is  well  marked.  There  are,  however,  other  instances  where  this  is 
not  the  case  and  where  it  is  therefore  very  difficult  to  know  whether  one 
is  dealing  with  a  compound  or  a  mixture.  We  shall  meet  with  many 
examples  of  this  later.  There  is,  however,  one  way  whereby  a  compound 
can  be  distinguished  from  a  mixture,  viz.,  by  ascertaining  whether  or  not 
the  substance,  prepared  in  different  ways,  has  a  constant  composition. 

PHENOMENA  ACCOMPANYING  THE  FORMATION  OR 
DECOMPOSITION  OF  A  COMPOUND. 

The  most  common  phenomenon  of  this  sort  is  an  elevation  or 
depression  of  temperature,  i.e.  an  evolution  or  absorption  of  heat 
(caloric  effect).  Sometimes  the  rise  of  temperature  is  so  great  that 
light  is  produced.  A  decomposition  or  a  combination  can  be  so 
violent  that  it  causes  an  explosion.  In  other  instances  electricity 
may  be  produced  by  chemical  action.  All  these  facts  may  be  com- 
prised in  this  statement:  Chemical  action  results  in  a  change  in  the 
energy-supply  of  the  reacting  substances. 

EXPLANATION    OF   THE    CONSTANT    COMPOSITION    OF 
COMPOUNDS.— ATOMIC  THEORY. 

21.  It  was  stated  that  constant  composition  is  the  distinctive 
characteristic  of  a  chemical  compound.  The  proportions  in  which 
elements  unite  to  form  a  certain  compound  are  always  the  same. 
This  Law  of  Constant  Composition  (definite  proportions)  was 
finally  established  by  PROUST  in  the  beginning  of  the  nineteenth 
century,  and  at  about  the  same  time  DALTON  offered  an  expla- 
nation of  it  which  is  still  accepted  and  may  be  considered  as  the 
foundation  of  theoretical  chemistry. 

This  explanation  involves  a  hypothesis  as  to  the  constitution 
of  matter.  It  is  possible  to  regard  matter  as  infinitely  divisible; 
according  to  human  conception  the  smallest  particle  that  can  really 
be  obtained  is  still  capable  of  division  into  an  infinite  number  of 
others.  However,  even  the  ancients  were  of  the  opinion  that  there 


28  INORGANIC  CHEMISTRY.  [§§21- 

must  be  somewhere  a  limit  to  the  divisibility  and  that  we  must 
finally  arrive  at  particles  incapable  of.  further  division,  the  atoms. 

In  the  fifth  century  B.C.  there  existed  a  school  of  philosophy,  that 
of  the  Eleatics  (so  called  from  the  city  of  ELBA),  whose  most  prominent 
representative  was  PARMENIDES.  He  taught  that  everything  that 
exists  cannot  be  otherwise  conceived  than  as  unchangeable;  trans- 
formation of  the  existent,  which  was  thought  to  have  never  originated 
and  to  be  at  the  same  time  unalterable,  was  held  by  them  to  be  incon- 
ceivable. These  theses  they  regarded  in  a  certain  sense  as  axioms, 
i.e.  statements  of  truths  which  are  accepted  without  proof.  Daily 
experience  teaches  one  nevertheless  that  transformation  does  occur  in 
that  which  exists,  a  fact  that  led  them  to  suppose  that  everything 
observed  by  men  is  merely  appearance. 

Three  theories  were  proposed  in  the  same  century  which  aim  to 
form  a  bridge  between  the  doctrine  of  the  unalterable  existent  and  the 
experience  that  points  toward  continuous  change.  These  theories 
originated  with  EMPEDOCLES,  ANAXAGORAS,  and  the  Atomists,  LEUCIPPUS 
and  DEMOCRITES.  The  immutability  of  the  existent  is  disposed  of  by 
ascribing  it  to  extremely  small  unchangeable  and  indestructible  particles ; 
every  change  is  thought  to  depend  on  the  movement  of  these  smallest 
integral  particles  toward  or  away  from  each  other.  EMPEDOCLBS  and 
ANAXAGORAS  assume  in  this  connection  an  infinite  divisibility;  the 
Atomists,  on  the  contrary,  regard  the  world  as  built  up  of  indivisible 
particles,  atoms,  all  of  which  consist  of  the  same  primordial  substance 
but  differ  in  form  and  size. 

Now  DALTON  has  used  this  conception  of  the  ancients  regarding 
the  atom  to  explain  the  fact  that  the  combining  weights  are  con- 
stant. The  atoms  of  the  various  elements,  he  assumes,  have  dif- 
ferent weights;  the  atoms  of  the  same  element  are  alike  in  weight. 
A  compound  of  two  elements  is  therefore  produced  by  the  associa- 
tion of  atoms  of  these  elements.  Such  a  combination  of  two  or 
more  atoms  is  called  a  molecule.  It  is  obvious  that  these  supposi- 
tions lead  directly  to  the  law  of  constant  proportions;  for,  if  copper 
oxide  is  formed  by  an  atom  of  copper  uniting  with  an  atom  of 
oxygen  to  make  a  molecule  of  copper  oxide,  its  composition  must, 
according  to  the  above  hypothesis,  be  constant. 

DALTON  deduced  another  conclusion  from  his  hypothesis,  and 
confirmed  the  same  experimentally.  He  observed  that  oxygen 
unites  not  only  with  one  very  definite  amount  of  nitrogen  oxide, 
but  also  with  twice  as  much,  not,  however,  with  any  intermediate 


22.]  THE  ATOMIC  WEIGHTS  OF  THE  ELEMENTS.  29 

amount,  He  also  showed  by  the  investigation  of  marsh-gas  and 
oiefiant  gas,  both  of  which  are  made  up  of  only  carbon  and  hydro- 
gen, that  the  former  contains  twice  as  much  hydrogen  to  a  certain 
weight  of  carbon  as  the  latter.  It  is  readily  seen  how  such  observa- 
tions can  be  explained  on  the  basis  of  the  atomic  theory;  in  one 
case  1  atom  of  carbon  is  in  combination  with  n  atoms  of  hydrogen; 
in  the  other  with  2n  atoms.  The  observations  of  DALTON  were 
subsequently  confirmed  and  extended,  especially  by  BERZELIUS. 
The  following  statement  is  therefore  now  accepted  as  a  law:  When 
two  elements  combine  to  form  more  than  one  compound,  the  different 
weights  of  the  one  element  which  unite  with  one  and  the  same  weight 
of  the  other  element  bear  a  simple  ratio  to  each  other.  This  is  the 
Law  of  Multiple  Proportions. 


THE  ATOMIC  WEIGHTS  OF  THE  ELEMENTS. 

22.  The  absolute  weight  of  the  atoms  is  only  approximately 
known  (see  §  35).  Nevertheless,  their  relative  weights,  i.e.,  the 
weights  of  the  atoms  of  the  various  elements,  when  that  of  a 
certain  element  is  arbitrarily  fixed,  have  been  determined  in  a 
variety  of  ways  (§§  208-210).  These  relative  weights  are  known 
as  atomic  weights. 

It  is  now  customary  to  take  the  atomic  weight  of  oxygen  as 
16.00.  The  atomic  weights  of  the  remaining  elements  then  have 
the  values  that  are  given  in  the  table  on  the  inside  of  the  back 
cover  of  this  volume.  The  acceptance  of  16  as  the  atomic  weight 
of  oxygen  has  a  historic  reason.  For  a  long  time  hydrogen  was 
taken  to  be  1 ;  it  was  believed  that  the  ratio  of  the  atomic  weights 
of  hydrogen  and  oxygen  was  1:16.  Inasmuch  as  the  atomic  weights 
of  most  elements  are  determined  from  the  composition  of  their  oxy- 
gen compounds,  the  basis  is  really  O  =  16  and  not  H=  1.  This  made 
no  difference,  so  long  as  the  proportion  H:0=l:16  was  con- 
sidered accurate.  Even  when  the  ratio  was  later  found  to  be 
a  different  one  (according  to  investigations  of  MOELEY  and  of 
W.  A.  NOTES  the  ratio  1:15.88  may  now  be  regarded  as  very 
accurately  determined),  it  was  still  the  simplest  plan  to  preserve 
O  =  16  as  the  basis,  since  a  change  would  necessitate  a  complete 
recalculation  of  all  the  atomic  weights,  and  this  necessity  would 


30  INORGANIC   CHEMISTRY.  [§§22- 

moreover  recur  as  often  as  a  new  refinement  of  methods  of  inves- 
tigation brought  about  a  change  in  the  ratio  H:O. 

A  few  years  ago  there  was  established  a  permanent  inter- 
national commission  whose  duty  it  should  be  to  revise  the  table 
of  atomic  weights  critically  every  year.  Those  values  are  accepted 
as  the  "  international  atomic  weights  "  which  appear  to  be  the 
most  probable  among  the  determinations  that  have  been  pub- 
lished. The  atomic  weights  in  the  table  are  carried  out  to  as 
many  decimal  places  as  may  be  accepted  with  certainty.  For 
many  purposes,  however,  it  is  sufficient  to  use  round  numbers, 
such  as  N=14,  Br  =  80,  etc. 

Besides  the  atomic  weights,  we  quite  frequently  use  equivalent 
weigfits.  These  are  the  weights  of  the  elements  which  combine  with  a 
unit  amount  of  a  certain  standard  element.  One  part  of  hydrogen 
combines,  for  instance,  with  35.5  parts  of  chlorine  and  with  8  parts  of 
oxygen.  These  amounts  of  hydrogen,  chlorine  and  oxygen  are  equiv- 
alent to  each  other.  The  atomic  weight  is  either  equal  to  the  equivalent 
weight  or  a  multiple  of  it. 

CHEMICAL  SYMBOLS  AND  FORMULAS. 

23.  The  relative,  or  atomic,  weights  are  expressed  by  symbols, 
that  were  introduced  by  BERZELIUS  and  are  of  great  convenience 
in  the  representation  of  compounds  and  the  formulation  of 
chemical  reactions.  The  symbols  whose  derivation  is  not  at 
once  apparent  are  taken  from  the  Latin  names  of  the  elements; 
e.g.,  Sb  from  stibium,  Au  from  aurum,  Cu  from  cuprum,  Hg  from 
hydrargyrum,  Pb  from  plumbum,  Sn  from  stannum,  Fe  from 
ferrum,  and  Ag  from  argentum. 

A  symbol  stands  not  only  for  the  element  concerned, 
but  also  for  the  relative  weight  of  an  atom  of  that  element.  If 
the  atomic  weight  of  copper  is  63.57  and  that  of  oxygen  16.00, 
the  symbol  Cu  indicates  63.57  parts  by  weight  of  copper,  the 
symbol  O  16.00  parts  by  weight  of  oxygen.  It  has  been  deter- 
mined that  in  copper  oxide  one  atom  of  copper  is  combined  with 
one  atom  of  oxygen;  copper  oxide  is  therefore  represented  by 
the  formula  CuO,  which  expresses,  first,  that  we  are  dealing 
with  a  compound  of  copper  and  oxygen,  and,  second,  that  1  atom 
(63.57  parts  by  weight)  of  copper  is  united  in  it  to  1  atom  (16.00 


24.]  STO1CH10METRICAL  CALCULATIONS.  31 

parts  by  weight)  of  oxygen.  Many  compounds  contain  several 
atoms  of  the  same  element.  This  is  indicated  by  placing  the 
proper  figure  to  the  right  of  and  below  the  symbol.  Sulphuric 
acid,  for  example,  contains  2  atoms  of  hydrogen  (H),  1  atom  of 
sulphur  (S)  and  4  atoms  of  oxygen  (0)  in  the  molecule.  Its 
formula  is,  therefore,  H2S04. 

Chemical  actions  can  be  very  simply  represented  by  the  use  of 
these  formulae;  thus,  the  decomposition  of  mercuric  oxide  into 
oxygen  and  mercury  by 

HgO  =  Hg+0; 

that  of  potassium  chlorate  into  oxygen  and  potassium  chloride  by 
KC103     =     KC1  +   30; 

Potass,  chlorate.  Potass,  chloride. 

the  generation  of  hydrogen  from  zinc  and  sulphuric  acid  by 


In  such  equations  the  same  atoms  and  the  same  number  of 
each  must  appear  on  both  sides,  in  accordance  with  the  principle 
of  the  Indestructibility  of  Matter. 

STOICHIOMETRICAL  CALCULATIONS. 

24.  If  the  formulae  of  the  compounds  are  known  —  the  means  of 
ascertaining  these  will  be  discussed  in  detail  later  —  and  the  atomic 
weights  of  the  elements  composing  them  also  known,  it  is  very  easy 
to  calculate  the  weights  that  enter  into  reaction  in  all  chemical 
changes.  A  couple  of  examples  may  serve  to  make  this  clear. 

1.  It  is  required  to  know  how  many  liters  of  oxygen  at  0°  and 
760  mm.  pressure  can  be  obtained  by  heating  1  kilogram  of  mercuric 
oxide. 

The  atomic  weight  of  mercury  is  200,  that  of  oxygen  is  16; 
mercuric  oxide,  HgO,  is,  therefore,  200  +  16.  Out  of  these  216 
parts  by  weight  of  mercuric  oxide  16  parts  of  oxygen  can  be  ob- 


32  INORGANIC  CHEMISTRY.  [§§24- 

tained  by  heating,  i.e.  from  1  kilo  (=1000  g.)  can  be   obtained 
=74.07  g.     Since  1  1.  oxygen  at  0°and  760  mm.  pressure 


weighs  1.4296  g.,  74.07  g.  occupy  a  volume  of       Q     =51.8  i 

j.  .TT^yo 


2.  How  much  water  can  be  formed  from  the  hydrogen  obtained 
by  the  interaction  of  1  kg.  zinc  and  the  corresponding  amount  of 
sulphuric  acid? 

The  reaction  of  zinc  and  sulphuric  acid  is  expressed  by  the 
equation 

Zn  +  H2SO4  =  ZnSO4  +  2H 

and  the  combustion  of  hydrogen  to  form  water  by  the  equation 

2H  +  O=H2O. 

From  these  equations  it  follows  that  the  hydrogen  formed  by 
the  action  of  1  atom  of  zinc  yields  1  molecule  of  water.  For  every 
atom  of  zinc  we  obtain,  therefore,  1  molecule  of  water.  The  atomic 
weight  of  zinc  is  65,  the  molecular  weight  of  water  18;  therefore 
65  parts  of  zinc  correspond  to  18  parts  of  water.  1  kg.  zinc  must 

.  ,  ,  1000X18 
yield  -  —  -  =  276.9  g. 

3.  How  many  grams  of  potassium  chlorate  are  necessary  to  pro- 
duce enough  oxygen  to  oxidize  500  g.  copper  to  copper  oxide? 

The  reactions  concerned  are 

KC1O3=KC1+3O    and    Cu+0=CuO. 

Hence  3  atoms  of  copper  can  be  oxidized  with  the  oxygen 
derived  from  1  molecule  of  potassium  chlorate.  For  every  3  atoms 
of  copper  1  molecule  of  potassium  chlorate  must  be  consumed. 
The  molecular  weight  of  the  latter  substance  is  39.10+35.6 
+  3X16  =  122.56;  the  atomic  weight  of  copper  is  63.57;  for  every 

122  56 
63.57  parts  of  copper  —  ^  —  =  40.85    g.    potassium    chlorate    are 

o 

500X40  85 

therefore  required.     Hence  500  g.  copper   require   -  = 

bo  .57 

321.5  g.  potassium  chlorate. 

In  most  chemical  computations  gram  molecules  are  employed, 
these  being  the  molecular  weights  of  the  substances  in  grams. 
The  abbreviation  mole  has  been  suggested  by  OSTWALD  for  this 


25.]  CHLORINE.  33 

term.     Thus  "  1  mole  "  copper  oxide  means  63.57+16.00  =  79.57 
grams  of  it. 

The  molecular  weight  in  milligrams  is  called  a  millimole.     In 
the  same  way  we  may  speak  of  a  kilomole,  etc. 


CHLORINE. 

25.  Chlorine  does  not  occur  free  in  nature,  since  it  acts  upon 
the  most  diverse  substances  at  ordinary  temperatures.  In  com- 
pounds, however,  it  occurs  extensively.  Common  table  salt  is  a 
compound  of  sodium  and  chlorine.  Various  other  metallic  chlor- 
ides are  also  met  with  in  nature. 

.     Chlorine  gas  can  be  obtained  by  the  direct  decomposition  of 
certain  chlorine  compounds;  thus: 

1.  By  the  electrolysis  of  hydrochloric,   or  muriatic,   acid  (i.e. 
a  solution  of  hydrogen  chloride,   HC1,  in  water).     Chlorine  is 
given  off  at  the  positive  pole  (anode),  hydrogen  at  the  negative 
pole  (cathode). 

The  indirect  decomposition  of  its  compounds  offers,  as  in  the 
case  of  hydrogen  (§  11),  the  most  practicable  methods  of  obtaining 
the  element.  They  are  all  based  on  the  oxidation  of  the  hydrogen 
of  hydrochloric  acid,  whereby  water  is  formed  and  chlorine  liber- 
ated. 

2.  Commercially,    as   well   as   in   the   laboratory,   manganese 
dioxide,  Mn02,  is  frequently  used  as  the  oxidizing  agent: 

Mn02  +  4HC1  =  MnCl2  +  2H20  +  2C1. 

It  is  very  often  convenient  to  generate  the  hydrochloric  acid 
from  salt  and  sulphuric  acid  in  the  same  vessel  with  the  manganese 
dioxide.  The  two  reactions  thus  proceed  simultaneously: 


I.  NaCl+H2S04 
II.  4HC1  +  MnO2  =  MnCl2  +  2H20  +  2C1. 

3.  Other  commonly  used  oxidizing  agents  are  chloride  of  lime 

and  potassium  dichromate;  e.g. 

K2Cr207  +  14HC1  =  2KC1  +  Cr2Cl6  +  7H20  +  6C1. 


34  INORGANIC  CHEMISTRY.  [§§25- 

4.  The  oxygen  of  the  air  can  also  serve  as  the  oxidizing  agent: 

2HC1+O=H2O+C12. 

For  this  purpose  a  mixture  of  60%  of  air  and  40%  of  hydrogen 
chloride  at  about  430°  is  passed  over  porous  bricks  which  are 
soaked  with  copper  sulphate  solution.  About  70%  of  the  hydro- 
gen chloride  is  converted  into  chlorine.  This  method,  which  is 
known  as  the  Deacon  process,  is  used  commercially.  The  copper 
sulphate  serves  as  a  catalyzer. 

The  progress  of  chemical  changes  is  often  modified  by  the 
mere  presence  of  a  substance  which  has  the  same  chemical  con> 
position  after  the  reaction  as  at  the  beginning.  Such  a  substance 
is  termed  a  catalyzer  and  the  action  which  it  exerts  is  called 
catalysis,  or  catalytic  action.  The  quantity  of  the  catalyzer 
necessary  to  exert  a  perceptible  influence  is  often  very  small, 
This  is  the  case,  for  example,  in  the  combination  of  hydrogen  and 
oxygen  in  the  presence  of  platinum  as  a  catalyzer  (§  13,  p.  16). 
A  minute  trace  of  platinum  sponge  brought  into  contact  with 
detonating-gas  accelerates  the  combination  to  such  a  rate  that 
the  reaction  takes  place  very  quickly  and  can  even  become 
explosive.  In  the  DEACON  process  a  small  quantity  of  copper 
sulphate  suffices  to  bring  into  reaction  unlimited  quantities  of 
hydrogen  chloride  and  oxygen.  At  the  temperature  of  ^430° 
there  is  practically  no  reaction  between  oxygen  and  hydrogen 
chloride  without  the  catalyzer.  That  there  must,  nevertheless, 
be  a  reaction,  although  a  very  slow  one,  can  be  demonstrated 
by  the  same  reasoning  as  in  §  13.  The  catalyzer  therefore 
does  not  cause  a  reaction,  but  only  accelerates  it.  OSTWALD 
compares  its  action  to  that  of  oil  on  the  axles  of  a  machine 
which  move  with  very  great  friction.  When  oiled,  the  ma- 
chine will  go  much  faster,  notwithstanding  that  the  force  of 
the  spring  (here  the  energy  of  the  chemical  reaction)  has  not 
changed.  A  further  point  in  the  analogy  is  that  the  oil  is  not 
consumed. 

In  most  cases  of  catalysis  it  can  be  proved  that  the  catalyzer 
takes  part  in  the  reaction  but  at  the  end  of  it  reappears  in  its 
original  condition.  In  the  platinum  catalysis  of  detonating-gas, 


27.]  CHLORINE.  35 

for  example,  the  metal  unites  with  the  oxygen,  whereupon  the 
resulting  compound  reacts  with  the  hydrogen,  giving  water  and 
metallic  platinum.  The  phenomenon  of  catalysis  is  universal. 
OSTWALD  thinks  it  probable  that  there  is  no  kind  of  chemical 
reaction  that  cannot  be  influenced  catalytically  and  that  there 
is  no  substance,  element,  or  compound,  which  cannot  act  as  a 
catalyzer. 

Catalyzers  may  accelerate  or  retard  reactions;  at  present, 
however,  much  more  is  known  of  the  first  than  of  the  second 
kind. 

26.  Physical  Properties. — Chlorine  is  yellowish-green  (hence  its 
name,  which  is  derived   from  jAc^pos,  greenish-yellow)  and  has 
a  disagreeable  odor.     Its  specific  gravity   is  2.45,  taking  air  as 
unity,  or  35.46,  based  on  O=16.     1  1.  chlorine  weighs,  therefore, 
3.208  g.  at  0°  and  760  mm.  pressure.     At  -34°  it  becomes  liquid 
under  ordinary  pressure;,  at  —102°  it  solidifies  and  crystallizes. 
Its  critical  temperature  is   146°.     Liquid  and  solid  chlorine  are 
yellow.      Chlorine  gas  dissolves  in  about  one-half  its  volume  of 
water.-    The  aqueous  solution  bears  the  name  " chlorine- water.". 
It  can,  therefore,  not  be  collected  over  water,  but  a  saturated  salt- 
solution  may  be  used,  in  which  it  is  only  slightly  soluble.     The  most 
convenient  way  to  fill  a  vessel  with  it  is  by  displacement  of  air,  the 
gas  being  conducted  to  the  bottom,  where  it  remains  and  drives  out 
the  air  above,  because  the  chlorine  is  denser. 

27.  Chemical  Properties. — Even  at  ordinary  temperatures,  chlo- 
rine combines  with  many  elements  and  acts  on  many  compounds. 
If  perfectly  pure  chlorine  is  mixed  with  an  equal  volume  of  hydro-, 
gen,  the  two  unite  in  direct  sunlight,  causing  an  explosion.     If  the 
chlorine  is  impure  or  the  sunlight  diffused,   combination  occurs 
slowly.     When  a  hydrogen  flame  is  introduced  into  chlorine  gas,  it 
continues  to  burn,  with  the  formation  of  hydrogen  chloride.    Many 
metals  combine  with  chlorine  with  the  evolution  of  light,  e.g.  cop- 
per (in  the  form  of  imitation  gold-leaf),  finely  powdered  antimony, 
molten  sodium,  etc.     The  precious  metals  are  in    general  quite 
resistive     to     chemical    action.     They  are,    however,    attacked 
by  chlorine  and  changed  to  chlorides,  i.e.,  chlorine  compounds. 
Gold,    for    instance,    dissolves    in    chlorine-water,    forming  gold 
chloride. 


36  INORGANIC    CHEMISTRY.  [§§27- 

Chlorine  also  unites  readily  with  many  non-metals,  e.g.  phos- 
phorus, which  burns  in  it  with  a  pale  flame  to  phosphorus 
chloride. 

The  tendency  of  chlorine  to  unite  with  hydrogen — its  so-called 
chemical  attraction,  or  affinity,  for  the  latter — is  so  strong  that 
chlorine  abstracts  the  hydrogen  from  many  hydrogen  compounds  in 
order  to  combine  with  it.  A  strip  of  paper  dipped  in  turpentine 
burns  with  a  sooty  flame  when  introduced  into  an  atmosphere  of 
chlorine;  the  chlorine  unites  with  the  hydrogen  of  the  turpentine 
and  sets  the  carbon  free.  A  burning  candle  continues  to  burn  in 
chlorine,  depositing  soot  (carbon)  and  forming  hydrogen  chloride. 
If  sulphuretted  hydrogen  gas,  H2S,  is  passed  into  chlorine-water, 
hydrochloric  acid  and  sulphur  are  formed. 

Water  is  also  decomposed  by  chlorine,  oxygen  being  liberated : 

2H20+2C12=4HC1  +  O2. 

This  reaction  takes  place  under  the  influence  of  sunlight,  but 
proceeds  very  slowly.  It  can  be  conveniently  demonstrated  as  in 
Fig.  14.  A  retort  is  filled  with  dilute  chlorine-water,  inverted 
and  exposed  to  the  sunlight.  After  a  few  days  a  bubble  of  gas 
collects  at  the  top  of  the  retort,  and.  on  investigation  with  a 
glowing  splinter,  it  is  found  to  be  oxygen. 


FIG.  14. — SLOW  DECOMPOSITION  OF  WATER  BY  CHLORINE. 


Upon  this  decomposition  of  water  depends  the  bleaching  and  disinfect' 
ing  action  of  chlorine  and  those  substances  which  generate  chlorine.  In 
bleaching,  the  coloring  matters — usually  of  an  organic  nature — are  oxi- 
dized by  oxygen  to  colorless  substances.  Bacteria  are  killed  by  oxida- 
tion. Ordinary  atmospheric  oxygen  does  not  produce  these  effects.  Lit- 
mus, for  instance,  which  is  rapidly  decolorized  in  moist  chlorine  gas,  is 


28.]  HYDROGEN  CHLORIDE.  37 

totally  unaffected  by  the  air.  The  particularly  energetic  action  of  the 
oxygen  that  is  produced  from  water  by  chlorine  is  explained  by  assuming 
that  it  exists  in  an  atomic  condition,  the  status  nascens,  regarding  which 
more  will  be  said  later  (§  38).  Perfectly  dry  chlorine  has  no  bleaching 
power. 

If  water  is  saturated  with  chlorine  at  0°,  crystals  are  deposited,  of  the 
composition  C12  +  8H2O,  chlorine  hydrate.  At  a  higher  temperature  these 
are  wholly  decomposed  into  chlorine  and  water. 


HYDROGEN    CHLORIDE,   HC1,   and  HYDROCHLORIC  ACID. 

28.  Hydrochloric  acid,  of  the  formula  HC1  (§  31),  is  a  gas, 
occurring  in  nature  in  the  free  state,  e.g.  in  the  gases  of  some 
volcanoes.  It  forms  an  important,  although  small,  part  of  the 
gastric  juice  of  man  and  other  animals. 

Some  of  its  methods  of  formation  have  been  already  given  (  §27), 
e.g.,  by  direct  synthesis  from  its  elements  under  the  influence  of 
light.  It  is  quite  remarkable,  however,  that  ultraviolet  rays 
decompose  hydrogen  chloride  even  at  ordinary  temperatures. 
We  saw  also  (I.  c.)  that  hydrogen  chloride  is  formed  by  the 
action  of  chlorine  on  hydrogen  compounds.  Moreover,  it  can 
also  result  from  the  action  of  hydrogen  on  some  chlorine  com- 
pounds, e.g.,  silver  chloride,  AgCl,  and  lead  chloride,  PbC^, 
when  heated  in  a  current  of  hydrogen,  yield  metal  and  hydro- 
chloric acid: 

AgCl+H=Ag  +  HCl. 

The  ordinary  method  of  preparation  is  by  the  action  of  a  chlorine 
compound  on  a  hydrogen  compound,  viz.,  that  of  salt  (sodium 
chloride)  on  concentrated  sulphuric  acid: 

NaCl  +  H2S04  =  NaHS04 + HC1. 

Sodium     Sulphuric 
chloride.         acid. 

This  method  is  employed   technically  as  well  as  in  the  labora- 
tory. 


38  INORGANIC  CHEMISTRY.  [§§28- 

The  above  reaction  takes  place  at  ordinary  temperatures.  If  the 
sulphuric  acid  is  to  be  completely  used  up,  i.e.  if  all  the  hydrogen  of  the 
sulphuric  acid  is  to  go  off  with  the  chlorine  of  the  salt  as  hydrochloric  acid, 
the  temperature  of  the  reaction  must  be  raised  (cf.  also  §  226) : 


2NaCl+  H2SO<  -Na^O,  +2HC1. 


29.  Physical  Properties.  —  Hydrogen  chloride  is  a  colorless  gas 
with  a  pungent  odor.  Its  critical  temperature  is  +52.3°;  the 
critical  pressure  86  atmospheres.  Liquid  hydrogen  chloride  boils 
at  -83.7°;  the  solid  melts  at  -111.1°.  Specific  gravity  of  the 
gas  =  1.2696  (air  =  l);  1  1.  HC1  at  0°  and  760  mm.  pressure 
weighs  1.6533  gr. 

For  obtaining  hydrogen  chloride  in  a  pure  state  MOISSAN  has  elabor- 
ated a  method  which  is  generally  applicable  to  gases,  since  low  tempera- 
tures are  easily  attainable  by  means  of  liquid  air.  The  freshiy  generated 
gases  contain  in  most  cases  moisture  and  other  impurities.  The  gases 
are  first  dried  by  being  passed  through  one  or  two  wash-bottles  placed 
in  a  bath  of  a  lower  temperature  than  —50°.  At  that  temperature  the 
tension  of  water  vapor  is  practically  zero.  The  gases  dried  in  this  way 
are  now  condensed  by  strong  cooling  to  the  solid  state.  Air  can  then  be 
pumped  out  of  the  vessel  If  the  temperature  is  now  allowed  to  rise,  the 
solid  mass  melts  first;  the  resulting  liquid,  when  vaporized,  gives  the 
perfectly  pure  gas. 

The  gas  fumes  strongly  in  the  air,  forming  a  cloud  with  the 
moisture  of  the  air.  It  is  very  soluble  in  water,  1  vol.  water 
at  0°  being  able  to  absorb  503  vols.  HC1  gas.  The  aqueous 
solution  of  the  gas  is  called  "  hydrochloric  acid,"  *  also  muriatic 
acid.  It  is  manufactured  commercially  on  a  large  scale  (§  226). 
Hydrochloric  acid  is  employed  almost  exclusively  in  the  form 
of  this  aqueous  solution.  A  solution  saturated  at  15°  contains 
42.9%  HC1  and  has  a  specific  gravity  1.212;  it  fumes  vigorously 
in  the  air.  The  ordinary  pure  "  concentrated"  or  "fuming" 
muriatic  acid  of  commerce  usually  has  a  specific  gravity  of  1.19 
and  contains  about  38%  HC1. 

*  The  gas  itself  is  often  called  ''hydrochloric  acid  gas"." 


30.J  HYDROGEN   CHLORIDE.  39 

Hydrogen  chloride  does  not  obey  the  law  of  HENRY  (§  9)  in 
its  behavior  towards  water,  for  its  solubility  in  this  liquid  is  not 
at  all  proportional  to  the  pressure.  The  larger  part  of  it  is  absorbed 
in  water  without  reference  to  the  pressure,  and  an  increase  of 
pressure  causes  only  a  small  increase  in  the  solubility.  Such  con- 
duct indicates  that  a  change  in  the  compound  has  occurred;  just 
what  this  change  consists  in  we  shall  soon  have  occasion  to  con- 
sider (§§  65,  66). 

The  saturated  solution  of  hydrogen  chloride  in  water  gives  off  HC1 
on  warming.  On  distilling  it  a  fraction  is  obtained  that  boils  con- 
stant at  110°  and  contains  8  mols.  H2O  to  1  mol.  HC1,  corresponding 
to  about  a  20%  solution  of  HC1.  A  solution  of  the  same  concentration 
and  boiling-point  results  from  distilling  a  more  dilute  hydrochloric 
acid,  enough  water  boiling  off  to  raise  the  concentration  to  the  above 
value. 

30.  The  chemical  properties  of  hydrogen  chloride  are  found  to 
be  quite  different  when  it  is  in  a  perfectly  dry  condition,  e.g.  con- 
densed to  a  liquid,  than  when  it  is  dissolved  in  water.  In  the 
former  case  it  does  not  act  on  metals  nor  change  the  color  of  blue 
litmus.  In  the  latter  case  just  the  contrary  is  true.  Zinc,  iron, 
and  other  metals,  when  dipped  in  the  aqueous  solution  of  hydrogen 
chloride,  are  vigorously  attacked,  hydrogen  being  given  off.  Blue 
litmus  is  turned  red  by  the  solution.  Moreover,  even  dilute  solu- 
tions taste  sour.  Now,  there  are  a  lot  of  substances  that  undergo 
a  similar  change  of  properties  when  they  are  brought  in  contact 
with  water,  and  whose  aqueous  solutions  possess  about  the  same 
properties  as  those  that  are  described  here  for  hydrochloric  acid. 
The  nature  of  this  change  will  be  discussed  later  on  (§  65).  It 
should  be  stated  here,  however,  that  these  substances  have  a 
common  name.  They  are  called  acids.  Acids  have  one  or  more 
hydrogen  atoms  that  can  be  replaced  by  metals.  The  compounds  of 
metals  that  are  formed  by  such  substitution  are  called  salts.  Salts 
can  result  not  only  from  the  direct  action  of  metals  on  acids,  but 
also  from  the  interaction  of  acids  and  bases.  The  term  "  bases " 
includes  compounds  of  the  general  type  MOH,  where  M  represents 
a  metal.  Most  of  them  have  an  alkaline  taste  and  turn  red  litmus 


40  INORGANIC  CHEMISTRY.  [§§30- 

blue.     When  sodium  is  dropped  into  water,  hydrogen  is  generated, 
and  a  base,  sodium  hydroxide,  is  formed: 

Na+H2O=NaOH  +  H. 

If  this  hydroxide  is  now  treated  with  hydrochloric  acid,  sodium 
chloride  and  water  are  produced: 

NaOH  +  HC1  =  NaCl + H20. 

If  we  indicate  an  acid  by  the  general  formula  AH  and  a  base 
by  MOH,  the  formation  of  salts  from  the  interaction  of  the  two 
may  be  represented  thus: 

MOH+HA=MA+H20. 

A  third  way  of  forming  salts  is  by  the  action  of  an  acid  upon 
a  metallic  oxide,  e.g. 

ZnO + H2SO4  =  ZnSO4 + H20. 

Zinc      Sulphuric          Zinc 
oxide.          acid.         sulphate. 

In  general,  the  bases  are  built  up  from  metals,  the  acids  from 
metalloids. 

When  hydrochloric  acid  is  added  to  a  solution  of  a  silver  salt,  fok' 
instance  to  silver  nitrate,  a  decomposition  of  this  salt  takes  place 
according  to  the  equation 

AgNO3 + HC1  =  HN03 + AgCl. 

Silver  nitrate.  Nitric  acid.     Silver 

chloride. 


The  silver  chloride  is  insoluble,  and  is  precipitated  as  a  white, 
curdy  mass.  In  this  reaction  the  hydrochloric  acid  has  liberated  the 
nitric  acid  from  its  salt.  It  is  also  possible  to  liberate  a  base  from 
a  salt  by  the  addition  of  another  base: 


AgN03  +  NaOH  =  AgOH  +  NaNO3. 

Silver  Sodium 

hydroxide.        nitrate. 

Such  reactions  are  called  single,  or  simple,  decompositions. 


31.]  COMPOSITION  OF  HYDROCHLORIC  ACID.  41 

Now  it  can  also  happen  that  two  salts  exchange  their  metals 
when  brought  together : 

Nad + AgN03 = AgCl + NaN03, 

Sodium 
chloride, 

so  that  two  other  salts  are  obtained.    Such  a  reaction  between 
salts  is  called  a  double  decomposition. 

We  shall  later  have  occasion  to  study  the  laws  governing  both 
of  these  decompositions. 


COMPOSITION  OF  HYDROCHLORIC  ACID.       LAWS  OF  GAY-LTJSSAC 
AND  AVOGADRO. 

31.  The  composition  of  hydrochloric  acid  is  determined  by  the 
following  experiments : 

(a)  When  strong  hydrochloric  acid  (a  more  than  23%  solution) 
is  subjected  to  electrolysis  in  a  suitable  apparatus  (see  below) 
it  is  observed  that  equal  volumes  of  hydrogen  and  chlorine  are 
evolved. 

(6)  Equal  volumes  of  chlorine  and  hydrogen  unite  to  form 
hydrochloric  acid  without  leaving  a  remainder  of  either  element. 
2  vols.  HC1  are  formed.  Since  the  weight  of  1  vol.  Cl  is  35.46 
(O=16),  hydrochloric  acid  must  consist  of  1  part  by  weight  of 
hydrogen  combined  with  35.46  parts  of  chlorine. 

In  the  electrolysis  of  hydrochloric  acid  sticks  of  charcoal  are  ordinarily 
used,  because  platinum,  the  substance  employed  in  most  other  electrolyses, 
is  attacked  by  chlorine.  The  apparatus  of  Fig.  11  is  also  impracticable, 
since  the  solubility  of  chlorine  in  water  increases  with  rising  pressure 
more  rapidly  than  that  of  hydrogen,  and  equal  volumes  of  both  gases  are 
therefore  not  obtained.  In  its  place  we  use  an  apparatus  suggested  by 
LOTHAR  MEYER  (Fig.  15),  by  which  the  compression  of  the  chlorine  by  a 
steadily  rising  column  of  liquid  is  avoided.  In  A  hydrochloric  acid  is 
electrolyzed  and  the  hydrogen  and  chlorine  are  collected  in  the  cylinders 
BB,  which  are  filled  with  a  saturated  sodium  chloride  solution.  The 
collected  gases  are  thus  under  diminished  pressure. 

The  combination  of  equal  volumes  of  chlorine  and  hydrogen  can  be 
carried  out  in  a  thick-walled  tube,  that  is  filled  with  the  gases  and  then 
exposed  for  a  day  to  diffused  sunlight.  Since  the  success  of  the  experi- 
ment requires  the  use  of  the  exact  proportions  of  chlorine  and  hydrogen 


42 


INORGANIC  CHEMISTRY. 


31- 


and  their  absolute  purity,  the  gas  mixture  is  prepared  by  electrolysis  in  the 
dark  and  exposed  to  the  action  of  light  immediately  after  the  tube  is  filled. 


FIG.  15. — ELECTROLYSIS  OF  HYDROCHLORIC  ACID. 

The  fact  that  hydrochloric  acid  gas  yields  a  volume  of  hydrogen  equal 
to  half  its  own  volume  can  also  be  shown  in  another  way.  When  perfectly 
dry  hydrogen  chloride  is  treated  with  sodium  amalgam — a  solution  of 
sodium  in  mercury — the  sodium  combines  with  the  chlorine,  setting 
hydrogen  free.  The  volume  of  the  latter  is  then  found  to  be  half  as  large 
as  that  of  the  hydrochloric  acid  taken. 

Hydrogen  and  chlorine  thus  unite  in  a  very  simple  ratio  by 
volume  (1:1),  and  the  volume  of  their  product  also  bears  a  very 
simple  ratio  to  that  of  the  components  (2:1:1).  In  discussing  the 
composition  of  water  (§  19)  we  already  remarked  that  oxygen  and 
hydrogen  combine  in  a  very  simple  ratio  by  volume,  viz.,  1:2. 
By  carrying  out  this  synthesis  at  a  temperature  above  100°, 
so  that  the  steam  is  not  condensed  to  water,  it  is  found,  further, 
that  the  volume  of  resulting  steam  bears  a  simple  ratio  to  the 
volumes  of  its  components,  viz.,  that  1  vol.  O  +  2  vols.  H  gives 
2  vols.  H2O. 

The  following  arrangement  serves  this  jwrpose  (Fig.  16).  The 
explosive  mixture  is  introduced  into  the  closed  arm  B  of  the  U-tube  over 
mercury.  B  is  surrounded  by  a  glass  jacket,  through  which  the  vapor  of 
boiling  amyl  alcohol  (generated  in  A),  whose  temperature  is  about  130°, 
is  passing.  This  vapor  is  condensed  in  C.  As  soon  as  the  gas  mixture 
has  reached  this  temperature  an  induction  spark  is  flashed  through, and  it  is 
found  that  the  volume  of  steam  formed  is  two-thirds  that  of  the  mixture. 


31. J 


COMPOSITION   OF  HYDROCHLORIC    ACID. 


FIG. 


-DETERMINATION  OF  THE  VOLUME  RELATIONS  BETWEEN  STEAM  AND 
ITS  COMPONENTS. 


What  was  found  above  to  be  true  for  hydrochloric  acid  and  for 
water  is  a  general  principle.  Gaseous  elements  combine  in  simple 
proportions  by  volume,  and  the  volume  of  the  products  formed — in  the 
gaseous  state — also  bears  a  simple  ratio  to  the  volumes  of  the  com- 
ponents. This  law  was  discovered  by  GAY-LUSSAC  in  1808. 

This  law,  together  with  the  atomic  theory  of  DALTON,  leads  to 
important  conclusions.  In  order  to  investigate  the  matter,  let  us 
assume  that  the  formula  of  hydrochloric  acid  is  HC1;  in  other 
words,  that  an  atom  of  hydrogen  is  in  combination  with  an  atom 
of  chlorine.  Since  one  volume  of  hydrogen  unites  with  one  volume 
of  chlorine  to  form  the  compound,  it  follows  from  the  above  as- 
sumption that  equal  volumes  of  chlorine  and  hydrogen  contain 
the  same  number  of  atoms. 

If  the  formula  were  otherwise,  e.g.  HnClOT,  the  numbers  of 
atoms  in  equal  volumes  of  hydrogen  and  chlorine  would  be  in  the 
ratio  of  n :  m. 

In  the  synthesis  of  water  2  vols.  of  hydrogen  and  1  vol.  of  oxy- 
gen yield  2  vols.  of  steam.  If  the  formula  of  water  be  H2nOp, 
the  numbers  of  atoms  in  equal  volumes  of  hydrogen  and  oxygen 
must  bear  to  each  other  the  ratio  n:p. 

Given,  therefore,  the  relative  numbers  of  atoms  in  equal  gas 


44  INORGANIC  CHEMISTRY.  [§  31- 

volumes  and  the  volume  ratio  in  which  the  gases  unite,  we  can 
determine  the  formula  of  the  resulting  compound. 

As  to  the  number  of  atoms  in  equal  gas  volumes,  there  was  at 
first  much  uncertainty.  Since  all  gases  behave  exactly  alike  to- 
wards changes  of  pressure  or  temperature,  it  was  reasonable  to 
suppose  that  the  number  should  be  alike  for  all  gases;  but  this  was 
soon  shown  to  be  incorrect.  In  the  synthesis  of  water  3  vols.  (2 
vols.  H+ 1  vol.  O)  give  2  vols.  of  steam;  hence  the  number  of  atoms 
per  unit  volume  must  be  different  for  steam  than  for  the  uncom- 
bined  elements.  However,  all  difficulties  were  overcome  by  a 
hypothesis,  which  AVOGADRO  enunciated  in  1811,  to  the  effect  that 
equal  volumes  of  all  gases  at  the  same  temperature  and  pressure  con- 
tain the  same  number  of  molecules. 

AVOGADRO  further  supposes  that  the  molecules  of  oxygen,  hy- 
drogen, chlorine,  and  other  elements  consist  of  two  atoms.  The 
union  of  hydrogen  and  chlorine  is  then  explained  thus :  Out  of  a 
molecule  of  each,  two  molecules  of  hydrochloric  acid  are  formed : 

H2  +  C12=2HC1. 

1  vol.  1  vol.     2  vols. 

The  total  number  of  molecules  thus  remains  the  same  after  the 
combination  and,  since  the  entire  volume  has  suffered  no  change 
either,  there  must'  be  just  as  many  molecules  present  in  each  of  the 
two  volumes  of  hydrochloric  acid  as  in  each  of  the  volumes  of 
hydrogen  and  chlorine. 

The  combination  of  hydrogen  and  oxygen  takes  place  thus: 

2H2+02=2H20. 

2  vols.  1  vol.      2  vols. 

Every  molecule  of  oxygen  has  split  up  into  its  two  atoms,  and 
each  of  these  unites  with  two  hydrogen  atoms.  The  number  of 
water  molecules  becomes  therefore  twice  as  great  as  that  of  the 
oxygen  molecules  and  equal  to  that  of  the  hydrogen  molecules; 
but,  since  the  volume  of  steam  is  also  double  that  of  oxygen,  there 
must  be  in  equal  volumes  just  as  many  water  molecules  as  oxygen 
molecules  and  hydrogen  molecules. 

32.  It  follows  from  the  above  that  AVOGADRO'S  hypothesis  is  of 
importance  hi  two  respects:  (1)  hi  furnishing  us  a  means  of  ascer- 


32.]  COMPOSITION  OF  HYDROCHLORIC  ACID.  45 

taining  the  relative  weights  of  the  molecules  of  gaseous  substances; 
(2)  in  putting  us  in  a  position  to  form  an  idea  of  how  many  atoms 
there  are  in  the  molecules. 

Let  us  examine  both  points  more  closely.  As  to  (1):  Since 
equal  volumes  of  gases  under  the  same  conditions  contain  the  same 
number  of  molecules,  the  ratio  of  the  weights  of  these  volumes  gives 
us  at  once  the  ratio  of  the  molecular  weights.  If  the  specific  gravity 
of  steam  is  9,  based  on  O=16,  and  that  of  hydrochloric  acid  is 
18.25,  the  ratio  of  the  molecular  weights  of  water  and  hydro- 
chloric acid  is  9: 18.25.  The  determination  of  the  specific  gravity 
of  gases  and  vapors,  the  vapor  density,  becomes  therefore  of  the 
greatest  importance  to  chemistry.  The  practical  method  of  pro- 
cedure is  described  in  ORG.  CHEM.,  §  12. 

For  the  determination  of  the  specific  gravity  of  gases,  see 
§212. 

As  to  (2) :  In  order  to  understand  how  AVOGADRO'S  hypothesis 
can  furnish  an  idea  of  the  number  of  atoms  which  the  molecules  oi 
elements  and  of  compounds  contain,  let  us  return  to  the  example 
of  the  synthesis  of  hydrochloric  acid.  1  vol.  hydrogen  unites 
with  1  vol.  chlorine  to  form  2  vols.  hydrochloric  acid.  According 
to  the  above  law  there  must  be  just  as  many  molecules  present 
in  the  two  volumes  of  hydrochloric  acid  as  there  were  molecules 
of  hydrogen  and  chlorine  together.  It  is  evident  that  this  is 
only  possible  in  case  the  molecules  of  hydrogen  and  of  chlorine 
divide  into  two  parts.  For,  if  the  chlorine  and  the  hydrogen 
molecules  consisted  of  only  one  atom  each,  the  volume  of  hydro- 
chloric acid  could  not,  in  accordance  with  AVOGADRO'S  law,  be 
double  that  of  each  of  its  elements,  but  would  have  to  be  equal 
to  it.  It  therefore  follows  that  an  even  number  of  atoms  must  be 
present  in  the  chlorine  and  in  the  hydrogen  molecules;  whether 
or  not  this  number  is  two,  as  AVOGADRO  assumed,  can  evidently 
not  yet  be  determined;  we  shall  therefore  represent  the  molecules 
of  hydrogen  and  of  chlorine  by  ~H.2x  and  C^j/.  From  the  synthesis 
of  water  the  same  conclusion  is  reached  in  regard  to  the  oxygen 
molecule:  2  vols.  hydrogen  unite  with  1  vol.  oxygen  to  form  2  vols. 
steam.  In  each  of  these  two  volumes  of  steam  there  must  be, 
according  to  AVOGADRO'S  law,  just  as  many  molecules  present  as 
in  the  one  volume  of  oxygen.  This  is  likewise  impossible  unless 
the  oxygen  molecule  splits  into  two  parts,  each  of  which  combines 


46  INORGANIC  CHEMISTRY.  [§§  33- 

with  a  molecule  of  hydrogen,  so  that  we  obtain  H2xOz  as  the  for- 
mula of  water  and  O23  as  that  of  the  oxygen  molecule. 

33.  The  formulae  for  hydrochloric  acid,  for  water  and  for  the 
molecules  of  hydrogen,  chlorine  and  oxygen  can  be  fully  established, 
if  the  values  x,  y  and  z  are  known.  These  can  be  ascertained  gen- 
erally in  the  following  way :  x  must  be  at  least  equal  to  1 ;  if  this  is 
the  case,  the  molecule  of  hydrogen  becomes  H2.  That  a  smaller 
number  of  atoms  is  impossible  is  shown  by  the  synthesis  of  hydro* 
chloric  acid.  The  vapor  densities  of  a  series  of  hydrogen  com- 
pounds, as  compared  with  that  of  hydrogen,  are  then  determined, 
from  which  we  can  find  their  molecular  weights,  based  on  the 
hydrogen  molecule  as  unity.  Thereupon  these  compounds  are 
analyzed  and  the  amount  of  hydrogen  calculated  that  is  repre- 
sented in  the  different  molecular  weights.  It  will  then  be  found 
that  in  no  case  is  the  amount  less  than  half  of  that  in  a  molecule 
of  hydrogen.  The  following  table  gives  some  examples: 

Substance.  Sp.  G.  (H=l).         Quantity  of  H. 

Hydrogen  chloride 18.25  0.5 

Hydrogen  bromide 40.5  0.5 

,Hydrogen  sulphide 17  1 

Ammonia  gas 8.5  1.5 

Methane 8  2 

Ethylene 14  2 

Water 9  0.5 

Since,  therefore,  no  compound  contains  less  than  half  a 
molecule  of  hydrogen,  the  atomic  weight  of  hydrogen  must  be 
half  its  molecular  weight,  i.e.  the  formula  of  the  hydrogen  mole- 
cule is  H2.  Similarly  it  is  found  that  the  oxygen  molecule  is 
O2,  that  of  chlorine  C12 ;  in  other  words,  that  x,  y,  and  z  are  all 
equal  to  1.  The  following  table  illustrates  the  case  of  oxygen: 

Substance.  Sp.  G.  (H  =  l).       Quantity  of  O. 

Oxygen 16  16 

Water 9  8 

Sulphur  dioxide 32  16 

Nitric  oxide 15  8 

Carbon  monoxide 14  8 

Carbon  dioxide .                               .  22  16 


34.]    DETERMINING  MOLECULAR  AND  ATOMIC  WEIGHTS.     47 

RULES  FOR  DETERMINING  MOLECULAR  AND  ATOMIC 
WEIGHTS. 

34.  When  the  atomic  weight  of  oxygen  is  taken  =16,  its 
molecular  weight  is  2  X 16  =32.  If  the  specific  gravity  of  another  gas 
based  on  oxygen=  16  is  a,  the  molecular  weight  of  this  gas  becomes 
2a.  The  following  rule  has  therefore  been  prescribed  for  the 
determination  of  the  molecular  weight.  Determine  the  vapor 
density  of  the  compound,  based  on  oxygen=16,  and  multiply  the 
result  by  2;  the  product  is  the  molecular  weight. 

For  determining  the  atomic  weight  the  following  holds  good, 
according  to  §  33 :  Determine  the  composition  of  molecular  amounts 
of  as  many  compounds  of  the  element  as  possible;  the  smallest  amount 
of  the  element  that  is  found  in  any  instance  is  the  atomic  weight. 

AVOGADRO'S  hypothesis  has  been  confirmed  from  a  physical  stand- 
point. It  is  at  present  one  of  the  principal  laws  of  chemistry  and  physics. 
Let  us  briefly  examine,  among  others,  the  physical  arguments  in  its  favor. 
The  molecules  of  bodies,  solids  as  well  as  liquids  and  gases,  are  in  con- 
stant motion,  the  intensity  of  which  increases  and  decreases  with  the 
temperature.  In  different  substances  at  the  same  temperature  there  must 
be  a  definite  relation  between  the  intensities  of  the  molecular  movements. 
This  relation  has  been  successfully  worked  out  from  the  theory  in  the  case 
of  gaseous  substances.  It  has  been  shown  that  in  all  gases  at  the  same  tem- 
perature the  mean  kinetic  energy  of  translation  of  a  molecule  is  the  same. 

The  pressure  which  a  gas  exerts  against  the  walls  of  the  vessel  is 
caused  by  the  impact  of  the  molecules.  We  will  call  the  number  of  mole- 
cules in  a  volume  of  the  gas  n,  the  mass  of  each  molecule  m  and  their 
mean  velocity  u.  It  is  then  clear  that  the  gas  pressure — the  above 
explanation  of  its  cause  being  accepted — must  be  proportional  to  n  and 
m.  Moreover  the  pressure  must  also  be  proportional  to  u2,  for  if  the 
velocity  were  increased  the  enclosing  walls  would  receive  more  impacts 
from  the  molecules  moving  to  and  fro,  and  every  impact  would  also 
become  stronger.  The  gas  pressure  p  is  therefore  proportional  to  the 

product  nmu2;    the  theory  says  that  p=%nmu2,  or  n=^T* 

In  this  expression  mu2  is  twice  the  kinetic  energy  of  translation  of 
molecules,  which  is  the  same  for  all  gases  at  the  same  temperature.  If 

then  p  is  made  the  same  for  the  different  gases,^^,  or  n,  the  number  of 
molecules  per  unit  volume,  must  be  the  same  for  all  gases. 

The  laws  of  BOYLE,  GAY-LUSSAC  and  AVOGADRO  (we  refer  to 
the  expansion  law  of  GAY-LUSSAC)  can  be  expressed  in  a  single 


48  INORGANIC  CHEMISTRY.  [§$  3l_ 

comprehensive  formula,  which  is  worthy  of  note  because  of  its 
frequent  use  in  physical  chemistry.  The  laws  of  BOYLE  and 
GAY-LUSSAC  are  represented  by  the  equation 

PV 
PV=RT,     or     ^-=#, 

in  which  P  is  the  pressure,  V  the  volume  and  T  the  absolute  tem- 
perature, of  the  gas  and  R  is  a  constant  which  depends  on  the 
quantity  and  the  nature  of  the  gas  under  consideration.  The 
value  of  R,  however,  becomes  the  same  for  all  gases,  if  molecular 
amounts  of  them  (One  mole  each)  are  taken.  For,  according  to 
AVOGADRO'S  law,  the  volume  of  one  mole  of  every  gas  is  the  same 
under  the  same  pressure  and  temperature.  In  the  'above  equa- 
tion, then,  V  is  constant  for  all  gases  and,  since  we  have  already 
made  P  and  T  the  same  in  each  case,  it  is  evident  that  R  must 
have  a  constant  value.  In  other  words,  if  we  deal  with  molecular 
amounts,  the  equation  PV=RT  becomes  a  general  expression 
of  the  laws  of  BOYLE,  GAY-LUSSAC  and  AVOGADRO. 

The  value  of  R  may  be  calculated  as  follows:  Let  us  consider 
1  mole  oxygen  at  0°  and  760  mm.  pressure.  Since  1  1.  oxygen 
under  these  conditions  weighs  1.4290  g.,  the  volume  V  of  1  molis 


If  a  correction  is  applied  because  oxygen  does  not  exactly  follow 
the  gas  laws  of  BOYLE  and  GAY-LUSSAC,  we  obtain  22412  c.c. 

The  pressure  of  760  mm.  mercury  corresponds  to  a  pressure 
of  1013.25  g.  per  sq.  cm.,  i.e.  P=  1013.25.  At  0°  the  absolute 
temperature  is  273°  (more  strictly  273.09°).  Substituting  these 
values  in  the  above  expression  for  R}  we  obtain 

PF_  1013.25X22412  _ 
=  T~=  273.09 

in  c.-g.  units.     If  the  pressure  is  expressed  in  millimeters  of  mer- 
cury, R  becomes 

760X22412 


273.09 


The  product  PV  also  represents  the  external  work  which  is 
'done  when  a  gas  under  constant  pressure  P  increases  its  volume 
by  V  (on  being  heated,  for  instance)  or  when  a  gas  being  generated 


35.]  THE  REALITY  OF  MOLECULES  AND  ATOMS.  49 

under  the  pressure  P  comes  to  occupy  a  volume  V.  For,  if  we  sup- 
pose that  the  gas  is  enclosed  in  a  cylinder  of  1  sq.  cm.  transverse 
section  having  at  one  of  its  ends  a  piston,  the  increase  of  the  volume 
must  cause  a  weight  P  to  move  through  V  cm.  One  calorie  (gram- 
calorie)  =4  1890  gram  centimeters.  If  this  is  substituted  in  the 


equation  PV  =  83155T,  the  latter  becomes  PV  =  T  ,  or,  very 

Q  -LOt/\J 

approximately,  PV  =  2T.  This  latter  form  also  is  a  common  one 
of  expressing  the  combined  gas  laws.  It  gives  the  external  work 
in  calories  that  is  done  when  1  gram  molecule  of  any  given  sub- 
stance is  converted  into  the  gaseous  state  at  the  absolute  tem- 
perature T. 

Since  1  gram  molecule  of  a  gas  has  a  volume  of  22.4  1.  at  0° 
and  760  mm.  pressure,  1  c.c.  under  these  same  conditions  contains 

—  ,  or  0.0446,  millimoles. 

THE  REALITY  OF  MOLECULES  AND  ATOMS  AND  THEIR 
ABSOLUTE  WEIGHT. 

35.  The  law  of  AVOGADRO  teaches  that  equal  volumes  of  all 
gases  contain  the  same  number  of  molecules.  If  we  take  a  gram 
molecule  oi  every  gas,  which,  as  we  just  saw,  has  a  volume  of 
22.41  1.,  it  follows  at  once  that  there  must  be  the  same  number 
of  molecules  in  every  case.  The  number  of  molecules  in  1  gram 
molecule  of  any  given  gas  is  thus  a  universal  constant;  it  is  often 
represented  by  N. 

The  determination  of  this  constant  N,  i.e.,  the  absolute 
number  of  molecules  contained  in  the  gram  molecule,  has  been 
worked  out  in  recent  years  by  several  widely  different  methods, 
all  of  which  have  yielded  approximately  the  same  result,  viz., 
70X1022. 

Back  in  1875  VAN  DER  WAALS  in  his  famous  treatise  on  the 
continuity  of  the  gaseous  and  the  liquid  states,  calculated  the 
value  of  N  to  be  between  40  and  90X1022,  which  is  of  the  same 
order  of  magnitude  as  the  present  more  accurate  value.  All  the 
methods  are  of  a  physical  character,  so  that  a  full  description  of 
them  is  inappropriate  here;  nevertheless,  to  show  the  diversity 
of  these  methods,  we  may  mention  that  the  above  value  of  N 
has  been  obtained  from  (1)  the  law  of  VAN  DER  WAALS;  (2)  the 
Browniai>  movement  (the  irregular  movement  which  solid  parti- 


50  INORGANIC   CHEMISTRY.  [§§35- 

cles,  having  the  dimensions  of  the  order  of  a  micron  (/*= 0.001  mm.) 
or  smaller,  exhibit  when  they  are  suspended  in  a  liquid.  It  has 
now  been  proved  that  these  movements  are  due  to  the  impacts 
of  the  molecules) ;  (3)  the  diffusion  velocity  of  dissolved  substances ; 
(4)  the  refraction  of  light  in  the  atmosphere,  causing  the  blue  color 
of  the  sky;  (5)  the  electric  charge  of  the  ions  (§266);  (6)  the  life 
period  of  radium  (§267);  (7)  the  energy  of  the  infra-red  spectrum. 
The  significance  of  these  investigations  for  all  departments  of 
natural  science  is  extraordinarily  great.  When  we  see  that  so 
many  wholly  independent  methods  lead  to  the  same  absolute 
number  of  molecules,  70  X 1022,  in  a  gram  molecule  there  is  no 
room  left  for  doubt  of  the  actual  existence  of  molecules.  So  long  as 
we  had  only  rough  and  discordant  approximations  of  this  number, 
we  could  accept  the  assumption  that  matter  is  built  out  of  mole- 
cules and  atoms  as  an  exceedingly  useful  hypothesis,  while  yet 
doubting  the  real  existence  of  atoms  and  molecules.  The  satis- 
factory proof  of  their  actual  existence  has  put  the  knowledge  of 
matter  in  general  on  a  secure  foundation. 

OZONE. 

36.  As  early  as  1785  VAN  MARUM  observed  that  when  an  elec- 
tric spark  passes  through  oxygen  a  peculiar  "  garlic-like  "  odor  is 
given  off,  and  a  bright  mercury  surface  is  at  once  made  dull. 
SCHONBEIN  investigated  this  phenomenon  more  carefully,  and 
found  that  it  is  due  to  the  formation  of  a  peculiar  substance,  which 
he  called  ozone.  This  proved  to  be  oxygen  existing  in  a  special 
condition.  The  fact  that  it  really  consists  of  nothing  but  oxygen 
is  shown  by  its  formation  from  perfectly  dry  oxygen  under  the 
influence  of  electric  discharges,  e.g.,  induction  sparks.  The  amount 
of  ozone  thus  formed  is  nevertheless  small.  It  is  greater  when 
silent  discharges  are  used.  As  ozone  is  formed  from  oxygen  by 
ultraviolet  light,  the  formation  of  ozone  by  silent  discharges  may 
be  caused  by  the  ultraviolet  light  accompanying  them.  This  is 
one  of  the  best  ways  of  obtaining  ozone,  although  the  maximum 
yield  is  only  5.6%.  However,  if  the  oxygen  is  cooled  by  liquid 
air  and  then  submitted  to  the  silent  discharge  at  a  pressure  of 
100  mm.  Hg,  it  is  wholly  converted  into  ozone. 

Fig.  17  represents  an  apparatus  constructed  by  BERTHELOT 
for  the  preparation  of  ozone  at  the  ordinary  pressure  and  tem- 
perature. The  wide  tube./,  together  with  the  supply-tube  d 


36.] 


OZONE. 


51 


FIG.  17. — PREPARATION  OF 
OZONE. 


and  the  exit-tube  e,  are  sunk  in  a  vessel  of  sulphuric  acid,  into 
which  the  pole  of  the  inductor  b  is  dipped.     The  other  wire  a  of 
the  latter  ends  in  a  tube  c,  which  is 
slipped    down    inside  /  and  is  almost 
entirely   filled.     The   silent   discharge 
between  the  two  bodies  of  sulphuric 
acid  thus  passes  through  a  thin  layer 
of  oxygen  and  has  a  powerful  ozoniz- 
ing effect. 

Ozone  is  formed  iu  many  reactions, 
such  as  the  slow  oxidation  of  moist 
phosphorus,  also  in  a  small  quantity, 
when  hydrogen  burns  in  an  atmosphere 
of  oxygen.  The  oxygen  that  is  ob- 
tained by  the  electrolysis  of  dilute 
sulphuric  acid  always  contains  it. 
Ozone  is  also  given  off  by  the  decom- 
position of  permanganic  acid  that  is  set  free  in  the  reaction  of 
potassium  permanganate  and  concentrated  sulphuric  acid  (cf.  also 
§52). 

When  oxygen  is  subjected  to  a  very  high  temperature  (e.  g.  flame 
temperature)  it  is  partially  converted  into  ozone,  and  the  more  so  the 
higher  the  temperature  (§103).  It  is  necessary,  however,  to  cool  down 
the  ozonized  gas  very  rapidly,  because  the  velocity  of  decomposition  of 
ozone  is  very  great,  especially  at  high  temperatures.  An  instantaneous 
cooling  can  be  accomplished  by  directing  the  flame  (of  hydrogen,  carbon 
monoxide,  acetylene  or  other  gas)  upon  the  surface  of  liquid  air,  which 
has  a  temperature  of  — 180° .  That  the  generation  of  ozone  has  no  con- 
nection with  the  combustion,  but  that  it  is  caused  only  by  the  high  tem- 
perature to  which  the  oxygen  is  raised  by  the  flame,  may  be  proved  by 
the  fact  that  an  incandescent  platinum  wire  or  NERNST  glower,  dipped 
in  liquid  air  also  generates  ozone.  The  formation  of  ozone  is  also  observed 
when  a  rapid  current  of  dry  air  or  oxygen  is  allowed  to  impinge  against 
a  hot  NERNST  glower.  When  the  air  contains  moisture  almost  no  ozone 
is  formed,  the  product  being  hydrogen  peroxide  (§54). 

Physical  Properties. — At  ordinary  temperatures  ozone  is  a  gas ; 
it  has  a  peculiar  odor,  which  is  one  of  the  most  delicate  tests 
for  its  presence.  One  part  of  ozone  can  still  be  detected  by  its 
odor  in  500,000  parts  of  air.  In  the  liquid  state  it  is  indigo-blue. 
Ozone  boils  under  normal  pressure  at  —119°. 


52  INORGANIC  CHEMISTRY.  [§§36- 

Chemical  Properties. — Ozone  is  characterized  above  all  by  its 
ability  to  oxidize  vigorously  at  ordinary  temperatures,  especially 
in  the  presence  of  moisture.  Phosphorus,  sulphur,  and  arsenic 
are  oxidized  to  phosphoric  acid,  sulphuric  acid,  and  arsenic  acid, 
respectively,  ammonia  to  nitric  acid,  and  silver  and  lead  to  per- 
oxides; e.g.,  the  metallic  surface  of  silver,  especially  when  heated 
to  above  240°,  become  blue  when  ozonized  air  is  directed  against 
it.  Iodine  is  deposited  by  ozone  from  a  solution  of  potassium 
iodide :  2KI  +  H  2O  +  O  =  2KOH  +  21. 

Organic  substances  are  strongly  oxidized  by  ozone,  hence  no 
apparatus  containing  it  should  have  connections  of  rubber.  Dye- 
stuff  solutions,  like  indigo  and  litmus,  are  decolorized  (by  oxida- 
tion) .  Ozone  effectively  destroys  micro-organisms,  and  is  there- 
fore used  successfully  in  the  sterilization  of  drinking-water. 

The  detection  of  ozone,  especially  in  quantities  too  small  to  be  recog- 
nized by  the  odor,  is  a  difficult  matter  because  several  other  oxidizing  sub- 
stances, such  as  chlorine  or  bromine  in  the  presence  of  water,  the  oxides  of 
nitrogen,  hydrogen  peroxide  and  still  others,  give  closely  analogous  reac- 
tions and  furthermore,  their  smell  at  high  dilutions  somewhat  resembles 
that  of  ozone ;  hence  it  becomes  necessary  to  first  prove  their  absence.  The 
tests  for  ozone  are  usually  executed  by  moistening  strips  of  filter-paper  with 
the  reagent  and  dipping  them  in  the  gas  containing  ozone.  The  reagents 
used  for  this  purpose  are  lead  sulphide  and  thallous  hydroxide.  The  strips 
of  paper  are  first  moistened  with  dilute  solutions  of  the  nitrates  of  these 
metals  and  then  exposed  to  hydrogen  sulphide  and  ammonia  fumes,  re- 
spectively. Lead  sulphide  is  oxidized  by  ozone  to  lead  sulphate,  thus  turn- 
ing from  black  to  white;  thallous  hydroxide,  which  is  white,  is  converted  to 
brown  thallic  hydroxide.  However,  these  changes  of  color  also  occur  with 
the  other  oxidizing  agents  mentioned.  A  strictly  characteristic  test  for 
ozone  is  the  violet  color  produced  with  an  acetic  acid  solution  of  tetramethyl- 
p-pf-  diamido-dipheriyl-methane  (an  organic  compound.)  Nitrogen  dioxide 
gives  a  straw-yellow  color,  chlorine  and  bromine  a  dark  blue,  while  hydrogen 
peroxide  produces  no  coloration  at  all. 

Ozone  is  stable  at  ordinary  temperatures,  but  is  easily  changed 
to  oxygen  on  heating.  It  is  slightly  soluble  in  water. 

37.  Formula  of  Ozone. — The  formula  of  ozone  has  been 
determined  by  LADENBURG  in  the  following  way.  A  glass 
globe  with  two  cocks  was  first  weighed  when  filled  with  pure 
oxygen  and  then  when  containing  ozonized  oxygen.  After 
reducing  both  weights  to  the  normal  temperature  and  pres- 
sure the  globe  in  the  latter  case  was  found  to  be  a  mg. 
heavier.  This  increase  of  weight  is  due  to  the  replacement 


38.]  HYDROGEN    PEROXIDE.  53 

of  a  certain  number  of  oxygen  molecules  by  the  same  number 
of  ozone  molecules. 

The  volume  that  the  ozone  occupies  in  the  gas  mixture  can  be 
determined  by  absorbing  it  in  turpentine.  Suppose  this  to  be 
v  c.c.,  when  reduced  to  normal  pressure  and  temperature.  The 
weight  of  this  v.  c.c.  ozone  can  be  represented  by  the  weight  of  an 
equal  volume  of  oxygen + a  mg.  and  must  be,  therefore,  (v  X  1.43 + a) 
mg.;  1.43  mg.  being  the  weight  of  1  c.c.  oxygen  at  normal  pressure 
and  temperature.  Hence  the  weight  g  of  1  c.c.  ozone  is 

0X1.43 -fa 

9=      —  - 

In  one  of  his  experiments  LADENBURG  found  a  =  16.3  mg.  and 

c\  f\n 

i>  =  26.0  c.c.,  hence  #=2.06  mg.     1  c.c.   ozone  thus  weighs  - 

1.45  times  as  much  as  an  equal  volume  of  oxygen,  or  very  nearly 
1J  times  as  much.  The  molecule  of  oxygen  being  O2,  that  of 
ozone  must  be  represented  by  63. 

In  an  oxidation  by  ozone  the  volume  of  the  ozoniferous  gas 
remains  unchanged.  Only  the  third  atom  in  Os  has  oxidizing 
power,  not  all  three  atoms  of  the  molecule. 

In  ozone  we  have  become  acquainted  with  oxygen  that  is  dif- 
ferent from  the  ordinary  kind.  This  phenomenon  is  also  seen  in 
other  elements;  it  is  called  allotropism. 

HYDROGEN   PEROXIDE,   H202. 

38.  This  compound  is  usually  prepared  by  treating  barium  per- 
oxide with  dilute  sulphuric  acid: 

Ba02  +  H2S04  -  BaS04  -f  H202. 

Barium  Insoluble, 

peroxide. 

In  a  very  concentrated  state  it  can  be  obtained  by  direct  distillation 
in  vacuo  of  a  mixture  of  sodium  peroxide  and  sulphuric  acid: 

NaA  +  H2S04=Na2S04+  H2O2. 

Hydrogen  peroxide  is  also  formed  in  many  other  ways;  e.g. 
together  with  ozone  (§  36)  in  the  slow  oxidation  of  phosphorus;  by 
the  combustion  of  hydrogen,  when  the  flame  is  cooled  by  a  piece 
of  ice.  The  formation  of  ozone  has  been  often  detected  when 
hydrogen  in  the  nascent  state  comes  in  contact  with  oxygen  mole- 
cules. We  suppose  that  in  the  moment  just  after  hydrogen  is  set 


54  INORGANIC  CHEMISTRY.  [§33. 

free,  its  atoms  have  not  yet  united  to  form  molecules,  so  that  the 
individual  atoms  possess  unusual  chemical  activity.  This  is  the 
general  conception  of  the  status  nascendi.  Thus  TRAUBE  has 
observed  the  following  instances  of  the  production  of  hydrogen 
peroxide:  Zinc  filings,  when  shaken  with  water  and  oxygen  or  air, 
give  hydrogen  peroxide,  since  the  zinc  and  the  water  generate  a 
small  quantity  of  hydrogen,  which  unites  with  the  oxygen.  Palla- 
dium-hydrogen behaves  likewise  when  brought  in  contact  with 
water  and  air.  In  this  case  it  is  the  hydrogen  released  from  the 
palladium  that  unites  with  the  oxygen.  Many  metals,  such  as 
copper,  lead  and  iron,  yield  hydrogen  peroxide  on  being  shaken 
with  air  and  dilute  sulphuric  acid,  for  the  same  reason  as  in  the 
case  of  zinc  and  water.  Finally,  the  peroxide  is  formed  in  the 
electrolysis  of  water,  when  a  current  of  air  or,  better,  oxygen  passes 
over  the  negative  electrode  (at  which  hydrogen  is  evolved). 

Hydrogen  peroxide  is  also  formed  at  very  high  temperatures 
from  steam  and  oxygen  (§  103);  just  as  in  the  formation  of  ozone 
under  the  same  conditions  (§  36), a  rapid  cooling  is  necessary  in 
this  case  also,  else  the  compound  decomposes.  The  formation 
of  hydrogen  peroxide  in  the  combustion  of  hydrogen  has  been 
shown  in  the  following  way:  A  hydrogen  flame  was  allowed  to 
burn  at  the  mouth  of  a  bulb  tube  containing  a  little  water.  By 
means  of  a  very  rapid  current  of  air  the  flame  was  blown  into 
the  bulb,  causing  a  very  sudden  cooling  of  the  mixture  of  steam 
and  air.  After  a  time  the  water  in  the  bulb  gave  the  tests  for 
hydrogen  peroxide.  As  a  further  analogy  to  the  case  of  ozone 
it  has  been  shown  that  the  formation  of  hydrogen  peroxide  has 
no  connection  with  the  combustion,  for  on  directing  a  fine  stream 
of  water  upon  an  incandescent  NERNST  glower  some  hydrogen 
peroxide  is  generated  in  the  water. 

Physical  Properties. — In  the  pure  anhydrous  condition  hydrogen 
peroxide  is  a  colorless,  slightly  viscid  liquid,  having  a  specific  gravity 
of  1.4584  at  0°,  based  on  water  at  4°.  (A  density  calculated  on 
this  basis  is  indicated  by  d£.)  It 'becomes  solid  at  a  low  tem- 
perature and  melts  at  —2°. 

Chemical  Properties. — Hydrogen  peroxide,  when  wholly  free 
from  impurities,  especially  from  suspended  particles  of  solid  mat- 
ter, is  rather  stable  and  can  be  distilled  in  vacuo;  when  impure, 
it  decomposes,  however,  into  water  and  oxygen,  as  it  also  does 
in  dilute  solution.  In  the  latter  state  it  is  more  stable  in  the 


§38.]  HYDROGEN    PEROXIDE.  55 

presence  of  traces  of  acid  than  in  the  presence  of  bases.  It  is  an 
interesting  fact  that  it  decomposes  rapidly  in  contact  with  powdered 
substances,  apparently  without  acting  upon  them.  Finely  divided 
silver,  gold,  platinum  (platinum  black),  and  especially  manganese 
dioxide  decompose  it  with  effervescence  (due  to  escaping  oxygen). 
Even  rough  surfaces  have  a  disturbing  effect;  BRUHL  observed, 
for  instance,  that  a  concentrated  solution  of  hydrogen  peroxide 
evolves  oxygen  when  poured  upon  ground  glass.  All  these  actions 
must  be  regarded  as  catalytic  accelerations  of  the  ordinarily  very 
slow  decomposition  of  hydrogen  peroxide.  The  effect  of  heat 
is  here,  as  elsewhere,  to  accelerate  the  reaction;  concentrated 
preparations,  when  warmed,  often  decompose  so  rapidly  as  to 
cause  an  explosion. 

The  oxidizing  action  of  hydrogen  peroxide  is  an  important 
chemical  property.  This  is  always  due  to  the  surrender  of  an 
oxygen  atom,  which  effects  the  oxidation,  while  water  remains. 
Lead  sulphide,  PbS,  is  oxidized  by  a  weak  solution  of  hydrogen 
peroxide'  to  lead  sulphate,  PbS04;  sulphuretted  hydrogen,  H2S, 
is  converted  into  water  and  free  sulphur.  Barium,  strontium  and 
calcium  hydroxides,  Ba(OH)2,  Sr(OH)2  and  Ca(OH)2,  are  pre- 
cipitated by  dilute  hydrogen  peroxide  from  their  solutions  as 
peroxides  of  the  general  formula  MO2<^aq.1  The  colorless  solu- 
tion of  titanium  dioxide  in  dilute  sulphuric  acid  is  turned 
orange-red  by  hydrogen  peroxide — lemon-yellow  by  traces  of  it 
— on  account  of  the  formation  of  yellow  trioxide,  TiOs.  This  is  a 
delicate  test  for  hydrogen  peroxide.  Other  tests  are  found  in  the 
following  oxidation  reactions:  Potassium  iodide  starch-paste  is  at 
once  turned  blue  by  hydrogen  peroxide  in  the  presence  of  a  little 
ferrous  sulphate,  FeSO4. 

The  ferrous  sulphate  carries  the  active  oxygen  of  the  hydrogen  per- 
oxide to  the  potassium  iodide.  As  a  result  two  atoms  of  iodine  are  set 
free,  the  ferrous  sulphate  being  oxidized  at  the  same  time.  According  to 
MANCHOT  a  higher  oxide  of  iron  is  formed  in  this  reaction. 

A  very  characteristic  reaction  is  this:  Chromic  acid  solution 
(H2CrO4),  when  treated  with  hydrogen  peroxide,  is  changed  to  a 
higher  oxide  (see  §  295)  which  is  blue  in  aqueous  solution  and 
may  be  taken  up  by  ether  if  shaken  with  the  latter.  This  test 
is,  however,  less  delicate  than  the  two  preceding  ones. 

1  Aq.  (aqua),  a  frequently  used  abbreviation  for  water  of  crystallization 
or  hydration.  **  _  * 


56  INORGANIC  CHEMISTRY.  [§§  38- 

A  third  group  of  chemical  effects  of  hydrogen  peroxide  depends 
on  its  reducing  power.  When  silver  oxide  is  introduced  into  a 
solution  of  hydrogen  peroxide,  a  vigorous  evolution  of  oxygen 
occurs,  water  and  metallic  silver  being  formed  at  the  same  time. 
Potassium  permanganate  solution  loses  its  color  when  mixed 
with  a  hydrogen  peroxide  solution  acidulated  by  sulphuric  acid, 
oxygen  being  given  off  rapidly: 

2KMnO4 + 3H2S04  +  5H2O2  =  K2S04 + 2MnS04 + 8H20  +  502. 

The  brown  peroxide  of  lead,  Pb02,  is  reduced  to  reddish-yellow 
lead  oxide,  PbO. 

Ozone  and  hydrogen  peroxide  yield  water  and  oxygen;  when 
dilute,  they  are,  however,  able  to  exist  side  by  side. 

There  is  a  test  for  hydrogen  peroxide,  depending  on  its  reducing 
power,  which  is  even  more  delicate  than  those  described  above.  A 
mixed  solution  of  ferric  chloride  and  red  prussiate  of  potash  has  a  red 
color.  On  the  addition  of  hydrogen  peroxide  Prussian  blue  is  precipitated. 
Traces  of  the  peroxide  turn  the  solution  green.  The  reaction  fails  in  the 
presence  of  free  acid. 

The  ability  of  so  powerfully  oxidizing  a  substance  as  hydrogen 
peroxide  to  act  also  as  a  reducing-agent  can  be  explained  as  follows : 
One  of  its  two  oxygen  atoms  must  be  loosely  joined  to  the  mole- 
cule, since  it  is  easily  given  up.  All  the  substances  which  are 
reduced  by  hydrogen  peroxide,  ajso  have  one  loosely  held  oxygen 
atom;  silver  oxide,  potassium/permanganate,  ozone  and  others 
give  up  their  oxygen  at  rather  low  temperatures.  It  is  therefore 
possible  that  the  mutual  attraction  of  the  oxygen  atoms,  which 
tends  to  make  them  form  oxygen  molecules,  is  stronger  than  the 
force  by  which  they  are  held  in  hydrogen  peroxide  on  the  one 
hand,  and  the  respective  oxygen  compound  on  the  other. 

Uses  of  Hydrogen  Peroxide. — The  colors  of  old  paintings  are  often 
restored  by  means  of  it.  The  darkening  of  them  is  due  in  many  cases  to 
the  transformation  of  white  lead  sulphate,  PbS04,  to  black  lead  sulphide. 
The  latter  is  readily  oxidized  by  hydrogen  peroxide  back  to  white  lead 
sulphate.  Hydrogen  peroxide  is  also  of  value  in  bleaching  ivory,  silk, 
feathers,  hair,  bristles  and  sponges.  It  is  also  important  in  analysis. 

For  therapeutic  purposes  at  30%  solution  of  hydrogen  peroxide  is  pre- 
pared by  MERCK  which  is  perfectly  pure  and  is  obtained  by  vacuum  dis  il- 
lation from  a  more  dilute  solution.  Before  use  it  is  strongly  diluted.  It 


40.]  DETERMINATION  OF  MOLECULAR  WEIGHT.  57 

has  the  advantage  of  not  being  subject  to  decompos'tion.  The  concentra- 
tion of  a  solution  of  hydrogen  peroxide  is  generally  expressed  in  the  volumes 
of  oxygen  that  it  can  evolve;  thus,  for  a  3%  solution  it  is  ten  volumes. 

39.  The  composition  of  hydrogen  peroxide  was  established  by 
THENARD  as  early  as  1818.  He  first  concentrated  it  in  a  vacuum 
and  then  introduced  a  weighed  amount  of  it,  enclosed  in  a  vial,  into 
a  graduated  barometer- tube  over  mercury.  The  vial  was  then 
broken  and  its  contents  decomposed  by  heating  the  tube  from 
without  or  allowing  finely  powdered  manganese  dioxide  to  rise  in 
the  tube.  It  was  thus  found  that  very  nearly  17  parts  of  hydrogen 
peroxide  by  weight  yield  8  parts  of  oxygen,  water  being  also 
formed.  One  atom  of  oxygen  (16  parts  by  weight)  is  therefore 
obtained  from  34  parts  of  hydrogen  peroxide,  the  remaining  18 
parts  forming  water;  in  other  words,  hydrogen  peroxide  is 
1  molecule  H2O  +  1  atom  0.  The  peroxide  therefore  contains  one 
atom  of  oxygen  to  every  hydrogen  atom.  Its  simplest  formula 
(the  so-called  empirical  formula)  is  then  HO.  Whether  this  also 
expresses  the  molecule  or  whether  the  latter  is  a  multiple  of  it, 
remains  to  be  determined  by  finding  the  molecular  weight,  inas- 
much as  every  compound  of  the  general  formula  (H0)n  possesses 
the  same  composition,  viz.,  16  parts  by  weight  of  oxygen  to  1  part 
of  hydrogen. 

On  account  of  the  instability  of  this  substance  its  vapor  density 
cannot  well  be  determined.  It  was  therefore  necessary,  in  finding 
its  molecular  weight,  to  follow  another  course,  which  is  based  on 
the  properties  of  dilute  solutions.  In  this  manner  the  molecule  of 
hydrogen  peroxide  was  found  to  possess  the  formula  H2O2.  The 
method  referred  to  is  explained  in  the  following  sections. 


MOLECULAR  WEIGHT    FROM   THE    MEASUREMENT   OF 

THE  DEPRESSION  OF  THE  FREEZING-POINT  AND 

ELEVATION  OF  THE  BOILING-POINT. 

40.  Certain  membranes  possess  the  peculiar  property  of  allow- 
ing a  solvent,  e.g.  water,  to  pass  through,  but  not  the  dissolved 
substances.  They  bear  the  name  "semi-permeable  membranes.''^ 
This  property  appears  to  depend  not  so  much  on  a  sort  of  sieve 
action  as  upon  the  ability  of  the  membrane  to  dissolve,  or  else  to 
absorb  or  loosely  combine  with,  the  solvent  on  one  side  and  release 


58  INORGANIC  CHEMISTRY.  [§§40- 

it  again  on  the  other,  while  the  dissolved  matter  remains  behind. 
One  of  the  ways  of  obtaining  a  semi-permeable  partition  is  by  dip- 
ping a  porous  cup — such  as  is  used  in  galvanic  cells — containing  a 
solution  of  yellow  prussiate  of  potash  into  a  solution  of  bhie  vitriol. 
A  thin  layer  of  copper  ferrocyanide  is  thus  formed  in  the  wall  of  the 
cup,  making  it  semi-permeable.  If  a  dilute  sugar  solution,  salt 
solution  or  the  like  be  poured  into  such  a  cup  and  the  cup  placed 
in  a  dish  of  water,  it  will  be  found  that  the  dissolved  substance  does 
not  diffuse  through  this  sort  of  a  partition.  The  water  goes 
through,  however,  for  if  the  cup  be  closed  with  a  perforated  stopper 
through  which  a  glass  tube  passes  and  then  dipped  deep  enough 
under  water  so  that  the  entire  cup  is  submerged,  the  water  will  be 
seen  to  rise  slowly  in  the  tube  till  it  reaches  a  definite  height  above 
the  level  outside. 

The  pressure  exerted  by  this  column  of  liquid  is  called  the 
osmotic  pressure  of  the  solution.  If  a  tight-fitting  piston  were 
inserted  in  the  cup,  the  force  which  one  would  have  to  exert  on  it 
to  prevent  the  infiltration  of  the  water  would  be  equal  to  the  pres- 
sure of  the  column  of  liquid,  for  the  water  continues  to  rise  in  the 
tube  till  the  pressure  of  the  column  prevents  the  entrance  of  any 
more. 

According  to  researches  of  VAN'T  HOFF  the  osmotic  pressure 
of  dilute  solutions,'  like  the  pressure  of  gases,  obeys  the  law  of 
BOYLE  and  the  expansion  law  of  GAY-LUSSAC.  If  the  pressure 
exerted  at  a  certain  temperature  by  a  kg.  of  a  gas  in  a  vessel  be 
p,  the  pressure  which  na  kg.  of  the  gas  at  the  same  temperature 
exerts  in  the  same  vessel  is  np.  The  concentration,  i.e.  density, 
of  the  gas  has  been  multiplied  n-fold. 

If  the  osmotic  pressure  of  a  solution  containing  a  per  cent  of  a 
substance  be  determined  and  found  to  be  p,  the  osmotic  pressure 
will  be  np,  if  an  na  per  cent  solution  of  the  same  temperature  be 
taken,  i.e.  if  the  concentration  be  n  times  as  great. 

An  investigation  of  the  pressures  which  a  gas  of  constant  volume 
exerts  at  the  absolute  temperatures  T\  and  T2  shows  that  these 
pressures  bear  to  each  other  the  ratio  TI  :  T2.  The  same  proportion 
is  observed  when  the  osmotic  pressure  of  a  solution  of  constant 
concentration  is  measured  at  the  same  absolute  temperatures  as 
above. 


41.]  DETERMINATION  OF  MOLECULAR  WEIGHT.  59 

41.  The  laws  of  osmotic  pressure  find  experimental  verification  in 
measurements  which  were  made  by  PFEFFER  previous  to  VAN'T  HOFF'S 
enunciation  of  the  laws,  PFEFFER  investigated  dilute  sugar  solutions  and 
used  an  apparatus  not  unlike  the  one  just  described. 

The  gas  laws  are  expressed  by  the  equation  (§  34) 

PV=RT, (1) 

in  which  P  represents  the  pressure,  V  the  volume,  and  T  the  absolute 
temperature  of  a  gas,  while  R  is  a  constant.  The  volume,  V,  is  inversely 
proportional  to  the  concentration,  according  to  the  above  definition; 

therefore  ^  may  be  substituted  for  V,  if  C  indicates  the  concentration. 
The  above  equation  then  becomes 

p 
C~RT' 


or,  at  a  constant  temperature, 


p 

= Const. 


This  equation  must  also  be  applicable  to  osmotic  pressure..  This  was 
really  the  case  in  PFEFFER' s  measurements  of  aqueous  sugar  solutions 
of  different  concentrations,  as  may  be  seen  from  the  following  brief 
table.  The  temperature  varied  between  13.5°  and  16.1°,  and  hence  was 
not  perfectly  constant: 

I 

1%  535  mm.  535 

2%  1016   "  508 

4%  2082    "  521 

6%  3075    "  513 

p 
The  differences  in  the  values  of  .^r  must  be  ascribed  to  the  variations 

of  temperature  and  the  unusual  experimental  difficulties  which  attend 
such  measurements. 

From  equation  (1)  it  also  follows,  when  V  (or  C)  is  a  constant,  that 

p 

-=  Const. 


60  INORGANIC  CHEMISTRY.  {§§  41- 

This  conclusion,  too,  was  demonstrated  experimentally  by  PFEFFER 
in  the  case  of  sugar  solution,  as  may  be  seen  from  the  following  table, 
A  one  per  cent  solution  was  used: 

P  T  £ 

P.  r 

510  287.15  1.78 

520.5  288.5  1.80 

544  305  1.78 

567  309  1.83 

VAN'T  HOFF  has  further  shown  that  the  numerical  value  of  the 
osmotic  pressure  is  the  same  as  that  of  the  gas  pressure;  that  is  to 
say,  when  a  definite  amount  of  a  substance  in  the  gaseous  state 
occupies  a  given  volume,  the  gas  pressure  which  it  exerts  is  just  as 
great  as  the  osmotic  pressure  which  would  be  produced  if  the  same 
amount  of  substance  were  dissolved  in  a  liquid  making  the  same 
volume  of  solution. 

The  measurements  of  PFEFFER  also  furnished  experimental  proof  of 
this.  He  found  that  a  1%  sugar  solution  at  7°  exerts  a  pressure  of  §  of  an 
atmosphere.  If  there  is  really  equality  between  osmotic  pressure  and  gas 
pressure,  or,  in  other  words,  if  the  law  of  AVOGADRO  for  gases  is  also 
applicable  to  dilute  solutions,  the  constant  R  of  the  equation  PV=RT 
must  have  the  same  value  for  solutions  as  for  gases.  P  in  the  above  case 
was  found  to  be  f  of  an  atmosphere,  or  |X  1033.6=689.0  g.  A  1% 
sugar  solution  contains  1  g.  sugar  in  100.6  c.c.  As  the  molecular 
weight  of  this  substance  is  342,  the  volume  V  which  contains  342  g.  is 
7=100.6X342.  T= 273  +  7° =280°.  Substituting  these  figures  in 

PV 
R  =  -jr,  we  have  #=84664.    The  close  agreement  of  the  two  values  of  R 

(compare  p.  48)  proves  the  equality  of  gaseous  and  osmotic  pressure. 

42.  It  follows  from  the  preceding  that  AVOGADRO 's  law  must 
also  hold  for  dilute  solutions.  Assuming  that  an  equal  number  of 
molecules  of  different  substances  are  dissolved  in  equal  volumes  at 
the  same  temperature,  we  know  from  the  equality  of  gas  pressure 
and  osmotic  pressure  that  the  various  substances  will  exert  the 
same  osmotic  pressure ;  inversely,  in  equal  volumes  of  solution  hav- 
ing the  same  temperature  and  osmotic  pressure  there  is  the  same 
number  of  molecules. 

This  is  a  very  important  extension  of  AVOGADRO 's  law.  We 
are  thus  able  not  only  to  compare  the  weights  of  equal  gas  volumes 
at  the  same  temperature  and  pressure  and  calculate  therefrom  the 
molecular  weight,  but  we  can  apply  the  same  principle  to  solutions, 


42.]  DETERMINATION  OF  MOLECULAR  WEIGHT.  61 

since  we  know  that  in  solutions  of  the  same  temperature  and  the 
same  osmotic  pressure  the  quantities  of  the  dissolved  substances 
contained  in  equal  volumes  of  solution  are  to  each  other  as  their 
molecular  weights. 

Just  as  it  is  possible  to  ascertain  the  molecular  weights  of 
gaseous  bodies  from  determinations  of  temperature,  pressure, 
weight  and  volume,  it  is  also  possible  to  find  those  of  substances 
in  dilute  solution  by  measuring  the  volume  of  liquid,  the  tempera- 
ture, the  quantity  dissolved  and  the  osmotic  pressure.  The 
molecular  weights  of  all  substances  that  dissolve  in  some  liquid  or 
otfrer  can  be  determined  in  this  way,  and,  since  the  number  of 
soluble  substances  is  very  large,  there  are  not  a  few  whose  molec- 
ular weights  were  first  determined  in  this  way. 

In  working  out  this  method,  however,  there  is  a  practical 
difficulty.  The  osmotic  pressure  is  very  hard  to  measure 
directly.  This  would  render  the  whole  method  of  little  value,  if  it 
were  not  for  the  fact  that  the  calculation  only  requires  that  it  be 
known  whether  two  solutions  have  the  same  osmotic  pressure,  not 
the  absolute  amount  of  the  latter;  the  law  of  AVOGADRO  simply 
requires  the  equality  of  volume,  of  temperature  and  of  pressure 
(osmotic  or  gas),  without  regard  for  the  absolute  value  of  these 
factors  (between  certain  limits).  Now,  it  is  easy  to  measure 
magnitudes  which  are  proportional  to  the  osmotic  pressure,  and 
from  which  it  may  be  seen  whether  equality  of  osmotic  pressure 
exists  or  not.  These  magnitudes  are  the  depression  of  the  freezing- 
point  and  the  elevation  of  the  boiling-point.  An  explanation  of 
these  terms  is  perhaps  necessary: — When  a  substance  is  dissolved 
in  a  liquid  the  maximum  tension  of  the  vapor  is  less  above  the  solu- 
tion than  above  the  pure  solvent  at  the  same  temperature,  for  the 
particles  of  the  dissolved  body  attract  the  molecules  of  the  solvent, 
hindering  the  formation  of  vapor  on  the  one  hand,  and,  on  the  other 
hand,  facilitating  the  return  of  vapor  molecules  into  the  liquid. 
This  lowering  of  the  vapor  pressure  necessarily  causes  a  depression 
of  the  freezing-point  and  an  elevation  of  the  boiling-point,  as  may 
be  proved  by  the  following  diagrams.  In  Fig.  18,  abc  represents 
the  vapor-pressure  curve  of  a  solvent  in  the  neighborhood  of  its 
freezing-point  6;  the  part  ab  gives  the  pressures  for  the  frozen 
matter;  the  part  be  for  the  liquid  solvent.  This  latter  part  is 
always  more  nearly  horizontal  than  the  former,  as  has  been  proved 


62 


INORGANIC  CHEMISTRY. 


[§§42^ 


both  experimentally  and  theoretically.  The  freezing-point  of  a 
liquid  is  that  temperature  at  which  the  solid  and  liquid  states  can 
exist  side  by  side  indefinitely.  This  condition  requires  that  the 
solid  and  the  liquid  substance  have  the  same  vapor  tension.  If, 
for  instance,  the  vapor  tension  of  the  solid  were  greater  than  that 
of  the  liquid,  we  should  have,  at  a  constant  temperature,  the  vapor 
given  off  from  the  solid  condensing  to  a  liquid  and  the  former  grad- 


FIG.  18. 


FIG.  19. 


ually  turning  into  the  latter.  Inversely,  if  the  vapor  tension  of  the 
solid  were  less  than  that  of  the  liquid,  the  entire  liquid  would, 
under  similar  conditions,  solidify. 

The  freezing-point  b  can  thus  be  regarded  as  the  intersection  of 
the  vapor-pressure  curves  ab  and  be  of  the  solid  and  the  liquid,  re- 
spectively. Let  us  now  consider  the  curve  &V  of  a  solution.  Its 
vapor  pressure  is  lower  than  that  of  the  pure  solvent,  so  its  inter- 
section with  the  curve  ab  must  lie  more  to  the  left,  that  is,  its 
freezing-point  is  lowered.  On  the  other  hand,  the  boiling-point 
of  a  solution  is  that  temperature  at  which  the  tension  of  its  vapor 
equals  one  atmosphere.  If  Od  in  Fig.  19  represents  this  tension, 
a  line  ddf  parallel  to  the  axis  of  abscissas  will  intersect  the  vapor- 
pressure  curve  ac  of  the  pure  solvent  at  a  lower  temperature,  than 
it  will  the  curve  a'c'  of  the  solution.  The  latter  must,  therefore, 
have  a  higher  boiling-point. 

43.  The  connection  between  these  magnitudes  and  the  osmotic 
pressure  will  be  better  understood  after  the  following  considerations: 

1.  Solutions  in  the  same  solvent,  separated  by  a  semi-permeable  parti- 
tion, can  only  be  in  equilibrium  when  they  are  isotonic,  i.e.  when  they  exert 
the  same  osmotic  pressure. 


43.] 


DETERMINATION  OF  MOLECULAR  WEIGHT. 


63 


Let  us  imagine  the  solutions  in  an  apparatus  consisting  of  two  cylinders 
that  are  connected  by  a  tube  containing  a  semi-permeable  partition.  In 
both  cylinders  the  level  of  liquid  is  kept  at  the  same  height  constantly  by 
adding  or  removing  some  from  time  to  time. 

The  solution  with  the  greater  osmotic  pressure  will  extract  solvent 
from  the  other,  for,  because  of  the  stronger  pressure  which  the  dissolved 
molecules  exert  upon  the  free  surface  of  the  liquid,  the  first  solution 
will  endeavor  to  increase  in  volume  at  the  expense  of  the  second. 
Equilibrium  will  be  established  so  soon  as  the  same  pressure  is  exerted 
by  the  dissolved  molecules  upon  the  unit  area  of  the  free  surfaces  of  the 
liquids  from  both  sides  of  the  semi-permeable  partition;  in  other  words, 
when  the  solutions  are  isotonic. 

2.  Isotonic  solutions  with  the  same  solvent  have  the  same  vapor  tension 
at  the  same  temperature. 


\ 

1 

A 

B 

H 

FIG.  20. 

The  proof  of  this  statement  lies  in  the  contradiction  to  which  the 
assumption  that  isotonic  solutions  have  unequal  vapor  tensions  leads. 
The  accompanying  diagram,  Fig.  20,  represents  a  closed  vessel,  that 
is  separated  by  the  semi-permeable  partition  HH  into  two  parts,  which 
contain  the  isotonic  solutions  A  and. B.  Near  the  top  the  two  parts 
are  connected  with  each  other.  Assuming  that  the  vapor  tension  of  A 
is  greater  than  that  of  B,  vapor  must  pass  out  of  A  and  condense  in  J5; 
the  result  is  that  A  becomes  more  concentrated,  B  more  dilute,  and 
they  are  no  longer  isotonic.  In  such  a  case,  according  to  the  first  prin- 
ciple, the  solvent  would  then  begin  to  pass  through  HH  from  B  to  A. 
The  assumption  of  perpetual  motion  which  is  thus  made  necessary  can 
only  be  avoided  by  supposing  that  the  vapor  tension  is  the  same. 
•  3.  Isotonic  solutions  with  the  same  solvent  have  the  same  freezing- 
point. 

Let  us  again  take  the  same  apparatus,  containing,  in  addition  to 
the  isotonic  solutions  A  and  B,  a  piece,  (7,  of  the  solvent  in  the  solid 
state  (Fig.  21).  Let  us  also  assume  that  A  and  C  have  the  same  vapor 
tension.  We  then  have,  according  to  definition  (see  §  42),  the  tern- 


64  INORGANIC  CHEMISTRY.  [§§  43,. 

perature  of  the  freezing-point  of  A.  However,  if  A  and  B  are  isotonic, 
they  have  the  same  vapor  tension.  B  will,  therefore,  have  the  same 
vapor  tension  as  C.  Hence  B  and  C  must  also  be  at  their  only  coexist- 
ence temperature,  the  freezing-point  of  B.  At  their  freezing-points 
A  and  B}  therefore,  have  the  same  temperature  as  C,  i.e.  they  possess 
the  same  freezing-point. 

4.  Isotonic  solutions  with  the  same  solvent  have  the  same  boiling-point. 

As  we  saw  in  §  42,  the  boiling-point  of  a  solution  is  that  temperature 
at  which  the  tension  of  its  vapor  equals  one  atmosphere.  Two  solutions 
with  a  common  solvent,  therefore,  have  the  same  vapor  tension  at  their 
boiling-point.  Now,  it  was  shown  above  that  solutions  having  the 
same  temperature  and  vapor  tension  are  isotonic.  If  these  solutions 
have  the  same  vapor  tension  (at  their  boiling-point)  and  are  isotonic, 
they  must  also  have  the  same  temperature. 

Since,  as  has  just  been  demonstrated,  isotonism  requires  like  freezing- 
points  and  boiling-points,  it  is  evident  the  depression  of  the  freezing- 
point  and  elevation  of  the  boiling-point  must  be  the  same  in  isotonic 
solutions  with  the  same  solvent. 

In  the  depression  of  the  freezing-point  and  the  elevation  of  the 
boiling-point  we  thus  have  a  means  of  deciding  whether  solutions 
are  isotonic.  Use  is  made  of  this  fact  for  the  determination  of 
molecular  weights  in  the  following  way:  The  freezing-point  of  a 
liquid,  e.g.  water,  acetic  acid,  phenol,  etc.,  is  first  determined. 
Thereupon  a  gram  molecule  of  a  substance  whose  molecular 
weight  is  known  is  dissolved  in  a  given  weight  (hence  also  in  a 
given  volume)  of  the  liquid.  A  depression  of  the  freezing-point 
is  observed.  This  depression  will  always  be  the  same,  no  matter 
what  the  substance  is  that  is  dissolved  in  the  liquid,  providing 
that  one  gram  molecule  is  dissolved  in  the  same  volume  of  liquid. 
The  depression  of  the  freezing-point  for  one  gram-molecule  of 
solute  is  thus  a  constant  for  the  solvent. 

Now  if  we  prepare  a  1%  solution  of  a  compound  whose 
molecular  weight,  M,  is  unknown  and  measure  the  freezing-point 
depression,  A,  we  have 

A  M= Constant. 

This  formula  is  also  applicable  to  the  elevation  of  the  boiling-point, 
as  can  be  readily  seen.  M  is  the  only  unknown  and  can  therefore 
be  calculated. 


44.]  BROMINE.  65 

When  water  is  used  as  the  solvent,  the  product  of  the  depression  A  of 
the  freezing-point  of  a  1%  solution  and  the  molecular  weight  M  has  been 
found  from  numerous  observations  to  be  19.  We  have  therefore  for  water 


For  hydrogen   peroxide,   the   depression  of  the  freezing-point  of  a 
3.3%  aqueous  solution  was  found  to  be  2.03°.     This  would  correspond  to 

2  03 

-^-3-  =0.615°  for  a  1%  solution;  hence  A  =0.615,  from  which  it  follows  that 

19 
the  molecular  weight  is 


Since  the  formula  HO  corresponds  to  a  molecular  weight  of  17,  H202 
to  one  of  34,  and  the  latter  number  is  the  nearer  to  the  molecular  weight 
found  by  experiment,  we  conclude  that  hydrogen  peroxide  has  the  doubled 
empirical  formula  H202. 

The  constants  for  the  freezing-point  depression  (molecular 
depression)  and  for  the  elevation  of  the  boiling-point  (molecular 
elevation)  of  some  compounds  that  are  well  adapted  for  these 
determinations  are  given  in  ORG.  CHEM.,  §  13. 

The  freezing-point  method  for  determining  molecular  weight 
is  called  the  cryoscopic  method,  while  the  boiling-point  method  is 
known  as  the  ebullioscopic  method.  Apparatuses  for  the  easy 
and  exact  determination  of  the  depression  of  the  freezing-point 
and  elevation  of  the  boiling-point  are  described  in  ORG.  CHEM., 
§§  14  and  15. 


BROMINE. 

44.  This  liquid  element  does  not  occur  free  upon  the  earth 
because  of  its  strong  tendency  to  form  compounds.  In  com- 
bination with  metals  it  is  found  in  the  salts  of  sea-water.  It  was 
discovered  in  the  latter  by  BALARD  in  1824.  Bromides  occur  in 
rather  large  amounts  in  the  so-called  Abraum-salze  of  the  Stassfurt 
salt-mines,  and  also  in  considerable  quantities  in  the  brines  of 
many  salt  wells,  notably  those  of  Michigan. 

In  the  neighborhood  of  Stassfurt,  Germany,  there  are  extensive  beds 
of  rock-salt  (halite).  Above  the  halite  are  found  layers  of  other  salts 
(called  "  Abraum-salze  "  because  they  have  to  be  removed  in  order  to  get 
at  the  halite).  These  salts  were  formerly  rejected  as  worthless,  but  they 


66  INORGANIC    CHEMISTRY.  [§§  44^ 

have  since  been  found  to  be  rich  in  potassium  salts,  bromides  and  other 
valuable  minerals,  so  that  the  "  waste  salts  "  of  former  days  are  now  the 
leading  source  of  many  commercially  and  scientifically  important  com- 
pounds. 

The  purification  of  these  Stassfurt  salts  is  accomplished  by 
solution  hi  water  and  partial  evaporation  of  the  latter.  Various 
substances  crystallize  out,  while  the  remaining  liquid  ("mother- 
liquor  ")  still  contains  the  most  soluble  salts,  among  which  is  mag- 
nesium bromide,  MgBr2.  From  this  mother  liquor  the  bromine  is 
obtained  by  the  use  of  chlorine,  which  sets  bromine  free  from 
bromides,  thus: 

MBr + Cl  =  MCI + Br .     (M  =  Metal.) 

The  process  employed  is  an  application  of  the  principle  of  the 
counter-current  (§15).  The  mother-liquor  is  allowed  to  flow  down 
through  a  tower  filled  with  round  stones,  so  that  the  exposed  sur- 
face of  the  liquid  is  greatly  enlarged.  A  current  of  chlorine  is 
passed  into  the  tower  from  below,  and,  as  it  rises,  the  gas  is  in  con- 
stant touch  with  the  bromide  liquor,  the  most  concentrated  gas 
being  in  contact  with  liquor  which  has  already  yielded  the  greater 
part  of  its  bromine,  so  that  practically  all  the  bromine  is  thus 
easily  obtained.  The  bromine  prepared  in  this  way  always  con- 
tains a  little  chlorine,  from  which  it  is  freed  by  distillation  over 
finely  powdered  bromide  of  potassium. 

Another  method,  common  in  the  United  States,  of  obtaining 
the  bromine  from  the  mother-liquor  is  by  distilling  the  latter 
with  manganese  dioxide  (or  potassium  chlorate)  and  sulphuric 
acid,  corresponding  to  the  method  of  making  chlorine  (§  25). 
Still  another  method  is  to  electrolyze  the  bromide  solution  and 
boil  off  the  bromine. 

The  bromine  thus  obtained  still  contains  a  little  water.  It  is 
dried  by  shaking  it  with  concentrated  sulphuric  acid  and  then  dis- 
tilling again. 

Physical  Properties. — Bromine  is  a  liquid  at  ordinary  tempera- 
tures; it  is  the  only  element,  excepting  mercury,  that  displays  this 
property.  It  solidifies  at  —7.3°  and  boils  at  59°.  It  is  dark  brown, 
and  is  transparent  only  in  thin  layers.  At  the  temperature  of 


45.]  HYDROBROMIC  ACID.  67 

liquid  hydrogen  (20.5°  absolute)  it  becomes  colorless;  MOISSAN 
showed  the  same  to  be  true  of  chlorine  and  fluorine  as  well.  It 
is  quite  volatile  at  ordinary  temperatures,  giving  off  brown  fumes 
of  an  extremely  irritating  and  disagreeable  odor,  whence  its  name 
(/3pSj^os  =  stench).  Sp.  g.  =3.1883  at  0°.  100  parts  of  water 
dissolve  3.5  parts  of  bromine.  The  addition  of  potassium  bromide 
to  the  water  increases  its  solubility  a  little.  Its  vapor  density  is 
79.96  (0  =  16). 

The  chemical  properties  of  bromine  are  completely  analogous  to 
those  of  chlorine,  but  the  action  of  the  former  is  less  energetic. 
While,  for  instance,  chlorine  combines  with  hydrogen  in  the  day- 
light at  ordinary  temperatures,  bromine  does  not.  Its  affinity  for 
many  elements  is,  however,  very  strong.  It  reacts  vigorously  with 
phosphorus;  and  powdered  arsenic  and  antimony  take  fire  when 
sprinkled  upon  bromine.  It  is  an  interesting  fact  that  of  the  two 
closely  related  alkali  metals,  potassium  and  sodium,  the  former 
reacts  vigorously  with  bromine,  while  the  latter  does  not  react  with 
it  at  all  at  ordinary  temperatures. 

The  bromine  molecule  consists  of  two  atoms ;  for,  since  its  vapor 
density  is  79.92  (see  above),  its  molecular  weight  must  be  159.84. 
Inasmuch  as  a  gram  molecule  of  no  one  of  the  very  numerous 
bromine  compounds  contains  less  than  79.92  g.  bromine,  but  often 
simple  multiples  of  this  quantity,  its  atomic  weight  is  taken  to 
be  79.92,  based  on  0=16.  The  molecule,  therefore,  contains 
159.84 


79.92 


2  atoms. 


HYDROGEN  BROMIDE,  or  HYDROBROMIC  ACID,  HBr. 

45.  This  gaseous  compound  can  be  obtained  by  direct  syn- 
thesis from  its  elements;  for  this  purpose  it  is  necessary  to  pass 
hydrogen,  together  with  bromine  vapor,  through  a  hot  tube  con- 
taining platinum  gauze.  This  is  the  most  practical  method  of 
manufacturing  it. 

Hydrobromic  acid  can  also  be  obtained  by  the  action  of  hydro- 
gen on  bromine  compounds.  Silver  bromide,  AgBr,  for  example, 
is  reduced  by  hydrogen  at  a  high  temperature  to  metallic  silver 
with  the  formation  of  hydrogen  bromide. 


68  INORGANIC  CHEMISTRY.  [§§  45- 

On  the  other  hand,  it  is  also  formed  by  the  action  of  bromine  on 
hydrogen  compounds.  For  this  purpose  numerous  organic  com- 
pounds can  be  used.  For  example,  bromine  reacts  with  naph- 
thalene, CioHg,  at  ordinary  temperatures  to  form  hydrogen  bromide, 
somewhat  impure,  however,  from  the  presence  of  organic  sub- 
stances. Hydrogen  bromide  is  also  produced,  together  with  free 
sulphur,  when  hydrogen  sulphide  'is  led  into  bromine  under  water: 

H2S+Br2  =  S+2HBr. 

Hydrobromic  acid  may  also  be  prepared  by  the  decomposition 
of  a  bromine  compound  with  a  hydrogen  compound,  phosphorus 
pentabromide,  PBr5,  and  water  being  employed  : 

PBr5  +  4H2O  =  5HBr 


Phosphoric 
acid. 

As  phosphoric  acid  is  not  volatile,  but  the  desired  substance  is,  the 
two  products  of  the  reaction  can  be  easily  separated. 

Physical  Properties.  —  At  ordinary  temperatures  hydrogen  bro- 
mide is  a  gas.  It  can  be  condensed,  by  cooling,  to  a  liquid  which 
boils  at  —64.9°  (under  738.2  mm.  pressure),  and,  by  still  farther 
cooling,  to  colorless  crystals,  which  melt  at  —88.5°.  It  has  a 
pungent  odor  and  a  sour  taste.  In  contact  with  moist  air  it  forms 
dense  clouds,  like  hydrochloric  acid  (§  29).  It  is  very  soluble  in 
water,  1  vol.  water  dissolving  about  600  vols.  at  10°;  its  solu- 
bility is  thus  even  greater  than  that  of  hydrochloric  acid. 

Chemical  Properties.  —  Here,  too,  the  acidic  nature  is  strongly 
displayed.  Various  metals,  such  as  zinc  and  magnesium,  are 
acted  upon  by  hydrobromic  acid,  forming  a  salt  and  free  hydrogen. 
The  most  of  its  salts  are  soluble  in  water;  silver  bromide,  however, 
is  insoluble  and  lead  bromide  difficultly  soluble. 

A  very  high  temperature  is  required  to  decompose  hydrogen 
bromide  into  its  elements. 

The  composition  of  hydrobromic  acid  can  be  determined  in  the 
same  way  as  that  of  hydrochloric  acid.  Since  its  vapor  density  is 
40.46,  it  has  a  molecular  weight  of  80.92.  The  atomic  weight  of 
bromine  being  79.92  (O=  16),  it  follows  that  the  formula  of  hydro- 
bromic acid  must  be  HBr.  Moreover,  the  dry  gas  can  be  decom- 


46.]  IODINE.  69 

posed  with  sodium  amalgam,  whereby  it  is  found  that  half  of  its 
volume  consists  of  hydrogen;  this  confirms  the  above  molecular 
formula. 

IODINE. 

46.  This  element,  a  crystalline  solid,  was  discovered  by  COUR- 
TOIS  in  1812,  but  its  elementary  nature  was  first  recognized  in  1815 
by  GAY-LUSSAC.  Like  chlorine  and  bromine,  it  does  not  occur  in 
the  free  state,  but  is  frequently  found  in  nature  in  combination 
with  some  metal.  An  important  source  of  iodine  compounds  is 
the  mother-liquor  (§  44)  remaining  in  the  purification  of  Chili 
saltpetre;  another,  the  ash  of  seaweeds,  known  in  Scotland  as 
kelp  and  in  Normandy  as  varec,  which  contains  iodides.  The 
extraction  of  the  iodine  is  accomplished  either  by  passing  chlorine 
into  the  solution  or  by  distilling  with  manganese  dioxide  and 
sulphuric  acid  in  the  same  manner  as  for  bromine  and  chlorine. 
The  commercial  iodine  is  purified  by  warming  it  gently  with  the 
addition  of  a  little  potassium  iodide,  the  iodine  subliming  in  the 
pure  state,  free  from  traces  of  chlorine  and  bromine  that  may 
have  been  present.  Finally,  it  is  dried  in  a  desiccator  over  sul- 
phuric acid. 

Physical  Properties. — Iodine  forms  tabular  crystals  of  a  dark 
gray  metallic  lustre.  Its  specific  gravity  is  4.948  at  17°.  It 
melts  at  114.2°,  and  boils  under  760  mm.  pressure  at  184.35°. 
Its  vapor  is  characterized  by  a  beautiful  dark-blue  color,  which 
gave  the  element  its  name  (io £tdrjs  =  violet).  In  water  it  is  only 
slightly  soluble — enough,  however,  to  color  the  water  yellow.  It 
dissolves  easily  in  a  solution  of  potassium  iodicje,  the  latter  being 
turned  brown.  In  various  other  liquids,  such  as  alcohol,  ether, 
carbon  disulphide  and  chloroform,  iodine  is  also  easily  soluble.  It 
is  a  peculiar  fact  that  the  alcoholic  and  the  ethereal  solutions  are 
brown,  while  the  solutions  in  carbori  disulphide  and  chloroform 
are  violet;  other  solvents,  e.g.  benzene,  give  solutions  of  an  inter-* 
mediate  color.  The  explanation  of  this  diversity  of  color  is 
that  in  the  brown  solutions  the  iodine  has  formed  a  compound 
with  the  solvent,  whereas  in  the  violet  solutions,  which  have  very 
nearly  the  same  color  as  iodine  vapor,  the  element  exists  in  the 
free  state. 


70  INORGANIC  CHEMISTRY.  [§46- 

This  conclusion  is  reached  in  various  ways;  one  is  from  the 
fact  that  the  addition  to  the  violet  iodine  solution  of  a  small 
amount  of  a  liquid  that  dissolves  iodine  with  a  brown  color  does 
not  alter  the  freezing-point  of  the  solution.  The  number  of 
molecules  free  to  move  has  thus  not  been  changed  by  this 
addition;  in  other  words,  the  iodine  has  united  with  the  added 
liquid. 

The  vapor  density  of  iodine  is  8.72  (air = 1)  at  about 
600°.  As  the  temperature  rises,  it  grows  steadily  smaller,  however. 
At  1500°  we  find  it  is  reduced  to  almost  half  of  what  it  is  at  600°. 
Later  (§  49)  we  shall  have  occasion  to  discuss  this  phenomenon, 
known  as  dissociation,  which  has  been  observed  with  many 
substances. 

47.  The  chemical  properties  of  iodine  resemble  very  strongly 
those  of  chlorine  and  bromine.  Its  affinity  for  other  elements  is 
in  general  weaker,  however,  than  that  of  the  two  halogens  men- 
tioned. It  combines  with  metals,  e.g.,  mercury,  directly  to  form 
salts  (iodides) .  A  characteristic  test  for  iodine  is  the  intense  blue 
coloration  which  it  imparts  to  starch  solution;  the  slightest  traces 
of  iodine  can  be  thus  detected.  The  blue  color  disappears  on 
boiling  and  reappears  on  cooling,  provided  the  boiling  was  not  too 
prolonged. 

The  blue  substance  formed  from  iodine  and  starch  is  not  a  com- 
pound, but  is  to  be  regarded  as  a  product  of  the  absorption  of  iodine  by 
starch;  for  the  quantity  of  iodine  taken  up  by  starch  from  the  solution 
in  KI  depends  largely  on  the  concentration  of  the  solution  and  does  not 
become  constant,  even  when  a  large  excess  of  solution  is  present.  It  is 
the  reaching  of  a  constant  ratio  that  characterizes  chemical  combination. 

The  molecule  of  iodine,  investigated  by  the  same  method  as  was 
employed  with  bromine,  is  found  to  consist  of  two  atoms  at  600°; 
hence  the  formula  is  I2.  Above  1500°  it  must  contain  only  one 
atom,  for  the  vapor  density  is  only  half  as  great. 


48.]  HYDROGEN  IODIDE.  71 


HYDROGEN  IODIDE,  or  HYDRIODIC  ACID,  HI. 

48.  This  compound  can  be  obtained  by  direct  synthesis  from 
its  elements,  and  that  is  really  the  best  method  for  preparing  it 
in  a  perfectly  pure  state.  For  this  purpose  hydrogen  and  iodine 
vapor  are  conducted  together  over  heated  platinum-black,  which 
accelerates  their  combination. 

Hydrogen  iodide  can  also  be  obtained  by  the  reaction  of  iodine 
with  hydrogen  compounds.  Organic  hydrogen  compounds  are 
preferable,  especially  colophonium  and  copaiva  oil.  This  method 
is  also  used  for  the  laboratory  preparation  of  the  gas,  but  the 
hydrogen  iodide  thus  obtained  is  more  or  less  adulterated  with 
organic  substances.  When  iodine  acts  on  hydrogen  sulphide  water 
hydrogen  iodide  is  formed,  sulphur  being  liberated  (§  45). 

As  an  example  of  the  action  of  hydrogen  on  an  iodine  compound, 
we  may  mention  the  reduction  of  silver  iodide  by  hydrogen,  from 
which  hydrogen  iodide  results. 

Finally,  the  action  of  an  iodide  on  a  hydrogen  compound  is 
illustrated  by  the  decomposition  of  a  phosphorus  iodide,  Pis  or 
PI5,  by  water.  As  was  explained  in  §  45,  it  is  possible  to  use 
phosphorus,  iodine  and  water.  This  method,  with  some  variation 
or  other,  is  the  preferable  one  for  the  preparation  of  hydriodic 
acid. 

According  to  GATTERMANN,  it  is  best  to  first  add  yellow  phosphorus 
(4  g.)  in  very  small  pieces  to  44  g.  iodine,  and  then  decompose  the  result- 
ing compound  with  a  little  water.  In  order  to  remove  the  free  iodine 
from  the  hydrogen  iodide  formed,  the  gas  is  allowed  to  pass  over  red 
phosphorus. 

The  decomposition  of  the  halogen  salt  by  sulphuric  acid  is  even  less 
available  for  the  preparation  of  hydriodic  than  for  hydrobromic  acid, 
since  the  former  is  more  easily  decomposed  by  sulphuric  acid  than  the 
latter. 

Physical  Properties. — Hydrogen  iodide  is  a  colorless  gas,  whose 
specific  gravity  is  62.94  (H  =  l).  It  fumes  strongly  when  exposed 
to  the  air,  and  possesses  an  acid  reaction  and  a  pungent  odor.  At 
0°  and  4  atmospheres  pressure  it  condenses  to  a  colorless  liquid, 
which  boils  at  —34.14°  under  a  pressure  of  730.4  mm.  The  melt- 
ing-point of  the  solid  is  —50.8°.  Hydrogen  iodide  is  very  soluble 


72  INORGANIC   CHEMISTRY.  [§§48- 

in  water;  1  vol.  H20  at  10°  dissolves  425  vols.  HI.  This  solution 
fumes  strongly,  and  turns  dark  brown  after  a  time,  because  of  the 
liberation  of  iodine. 

Chemical  Properties. — Hydrogen  iodide  has  all  the  charac- 
teristics of  an  acid.  With  metals  it  forms  salts  (iodides),  hydrogen 
being  given  off.  These  are  almost  all  soluble,  with  the  exception 
of  the  iodides  of  silver,  and  mercury.  Lead  iodide  is  slightly  soluble 
at  ordinary  temperatures.  In  addition  to  its  acidic  character, 
hydriodic  acid  possesses  another  property,  which  is  not  found  in 
hydrochloric  and  hydrobromic  acids.  Since  it  splits  up  readily 
into  hydrogen  and  iodine,  it  can  act  as  a  strong  reducing-agent, 
especially  at  high  temperatures.  It  has  already  been  remarked 
that  the  aqueous  solution  of  the  gas  turns  brown,  iodine  being 
set  free  by  the  oxidizing  action  of  the  air;  this  change  is  greatly 
aided  by  the  influence  of  light.  At  a  high  temperature  hydrogen 
iodide  is  decomposed  into  H2  and  12,  as  is  shown  by  the  appearance 
of  the  violet  iodine  vapor.  In  organic  chemistry,  particularly, 
frequent  use  is  made  of  the  reducing  power  of  this  acid. 

Formula  of  Hydriodic  Acid. — The  vapor  density  of  this  sub- 
stance has  been  found  to  be  62.94.  Its  molecular  weight  is  there- 
fore 125.88.  The  atomic  weight  of  iodine  being  126.92  (0=16), 
it  is  seen  that  the  molecular  weight  corresponds  very  closely  to  the 
formula  HI  =127. 92,  and  no  other  formula  is  possible. 

DISSOCIATION. 

49.  When  hydrogen  iodide  is  subjected  to  a  slow  increase  of 
temperature,  it  commences  to  decompose  at  a  definite  temperature 
slightly  above  180°  into  hydrogen  and  iodine  vapor.  As  the  heat- 
ing continues,  the  decomposition  grows  gradually  greater  till  a 
point  is  finally  reached  when  the  gas  mixture  contains  only  the 
individual  elements.  If  the  mixture  is  then  slowly  cooled,  the 
same  stages  are  passed  through  in  inverse  order,  so  that  the  degree 
of  decomposition  at  any  one  temperature  is  found  to  be  the  same, 
no  matter  whether  the  temperature  was  approached  from  above 
or  below,  it  being  only  necessary  that  the  particular  temperature 
should  be  maintained  for  a  sufficient  length  of  time  in  both  cases. 

The  phenomenon  just  described  is  to  be  observed  with  a  great 


49.]  DISSOCIATION".  73 

many  substances.  It  is  called  dissociation,  and  was  first  studied  in 
1857  by  H.  SAINTE-CLAIRE  DEVILLE. 

The  degree  of  decomposition  of  hydrogen  iodide  is,  as  we  have 
seen,  always  the  same  for  a  definite  temperature.  It  necessarily 
follows  from  this  that  if  one  starts  with  the  uncombined  elements, 
hydrogen  and  iodine,  and  heats  them  together  long  enough  at  a 
certain  temperature,  a  gas  mixture  of  exactly  the  same  compo- 
sition will  be  formed  as  that  resulting  from  the  decomposition  of  the 
hydrogen  iodide  at  the  same  temperature.  This  is  confirmed  by 
experiment;  e.g.  equivalent  amounts  of  iodine  and  hydrogen  were 
heated  in  a  sealed  vessel  by  exposing  it  to  the  vapor  of  boiling  sul- 
phur (445°).  The  amount  of  hydrogen  iodide  finally  produced 
was  79.0%,  i.e.  21.0%  of  the  gas  mixture  remained  uncombined. 
Again,  when  a  like  vessel  filled  with  hydrogen  iodide  was  heated  to 
the  same  temperature,  it  was  found  that  21.5%  had  decomposed 
—  a  figure  very  close  to  that  obtained  in  the  preceding  experiment. 

Such  reactions,  which  lead  to  the  same  result,  no  matter 
whether  we  start  with  the  one  set  of  substances  (H2  +  I2)  or  with 
the  other  (2HI),  are  called  reversible  reactions.  When  the  final 
stage  is  reached,  the  sets,  or  "  systems,"  are  said  to  be  in  equi- 
librium with  each  other. 

If  we  have  a  system  of  substances  A+B+  .  .  .  ,  which  is  par- 
tially changed  into  another  system  P+Q+  .  .  .  ,  the  equilibrium 
between  the  two  systems  is  expressed  by  the  sign  <=±;  thus: 


We  saw  that,  in  the  preparation  of  hydrogen  iodide,  platinum  black 
is  used  because  it  accelerates  the  combination.  Neither  this  nor  any 
other  catalyzer,  however,  changes  the  proportional  extent  to  which 
combination  takes  place.  For  instance,  experiments  have  shown  that 
at  350°  18.6%  of  the  hydriodic  acid  is  decomposed  when  no  platinum- 
black  is  present,  and  that  in  the  presence  of  this  catalyzer  the  decom- 
position reaches  19%;  these  two  figures  are  alike  within  the  limits  of 
experimental  error.  There  is  a  theoretical  reason  why  this  must  be  so. 
If  the  catalyzer  influenced  the  equilibrium,  we  could  realize  combination 
and  decomposition  by  alternately  adding  and  removing  the  catalyzer. 
Under  constant  conditions  of  temperature  the  system  would  absorb 
energy  in  one  instance  and  evolve  it  in  the  other.  The  energy  obtained 
in  the  latter  case  could  be  used  to  do  work.  Work  would  thus  be  gained 


74  INORGANIC  CHEMISTRY.  [§§49- 

from  a  system  remaining  at  constant  temperature,  but  according  to  the 
principles  of  thermodynamics  this  is  impossible,  since  the  production  of 
work  is  always  accompanied  by  a  fall  of  temperature. 

The  question  now  arises,  how  such  an  equilibrium  comes  about, 
and  why  the  decomposition  of  a  compound  that  begins  at  a  cer- 
tain temperature  does  not  complete  itself.  To  this  the  kinetic 
theory  of  gases  furnishes  a  satisfactory  answer.  According  to  this 
theory  the  molecules  of  gases  are  constantly  in  motion.  While 
a  constant  mean  velocity  of  the  molecules  may  be  assumed  to  exist 
for  every  temperature,  the  velocities  of  the  individual  molecules 
must  be  considerably  different,  because  of  their  very  frequent 
collisions  with  each  other.  The  atoms  of  a  molecule  must  also  be 
supposed  to  be  capable  of  changing  their  respective  positions,  for 
the  repeated  collisions  of  the  molecules  displace  the  atoms  from  their 
positions  of  equilibrium.  These  movements  of  the  atoms  are  the 
more  violent  the  greater  the  velocity  of  the  molecules.  It  is  easy 
to  conceive  that  they  may  at  last  become  so  violent  as  to  throw 
the  atoms  out  of  their  sphere  of  mutual  attraction.  The  molecule 
is  thus  broken  up.  In  a  body  of  gas  at  a  definite  temperature  this 
will>  however,  only  occur  in  those  molecules  whose  velocity  is 
above  certain  limits;  hence  we  see  a  reason  for  partial  decomposi- 
tion. The  explanation  of  partial  combination  is  exactly  analogous. 
The  atoms  set  free  from  the  molecules  of  the  elements  enter  into 
the  spheres  of  mutual  attraction,  and  if  their  velocities  are  not 
great  enough  to  resist  the  attraction,  the  different  atoms  unite. 

In  the  case  of  the  formation  and  decomposition  of  hydrogen  iodide 
this  can  be  conceived  as  follows:  Two  HI  molecules  meet  in  such  a 
state  of  atomic  movement  that  the  H  atoms  enter  the  spheres  of  attrac- 
tion of  each  other,  and  the  I  atoms  likewise,  so  that  H2  and  I2  are  formed. 
On  the  other  hand,  these  molecules  H2  and  I2  may  again  meet  in  such  a 
way  that  each  H  atom  enters  the  sphere  of  attraction  of  One  of  the 
I  atoms,  whereupon  two  HI  molecules  are  formed. 

According  to  the  above  the  state  of  equilibrium  is  to  be  ac- 
counted for  by  supposing  that  in  the  unit  of  time  just  as  much 
passes  from  the  one  system  into  the  other  as  vice  versa.  Until 
the  state  of  equilibrium  is  reached,  the  amounts  which  pass  from 
one  system  into  the  other  in  the  unit  of  time  are  unlike. 


50.]  DISSOCIATION.  75 

In  order  to  define  more  clearly  this  condition  of  equilibrium  we 
must  introduce  the  concept  of  reaction  velocity.  By  this  term 
we  understand  the  number  of  moles  transformed  from  one  system 
into  the  other  in  the  unit  of  time.  Suppose  that  in  the  unit  of 
volume  (one  liter,  for  instance)  there  are  a  moles  of  a  substance 
A,  which  can  undergo  a  chemical  change  into  a  substance  B.  If 

in  the  unit  of  time  (say  one  minute)  —  moles  of  A  are  converted 

into  J5,  the  reaction  velocity  S  will  be  expressed  by  —  .    Incase, 

n 

there  are  originally  only  Ja  moles  of  A  per  liter,  experiments 
have  shown  that  the  number  of  moles  converted  per  minute  is 

zr—  .    We  thus  perceive  that  the  reaction  velocity  is  proportional 

£  11 

to  the  number  of  gram-molecules  per  liter,  or,  in  other  words,  to 
the  concentration.  This  is  a  principle  of  very  wide  application4, 
it  is  ordinarily  called  the  law  of  chemical  mass  action.  It  finds 
a  general  expression  in  the  equation 


in  which  K  is  a  constant  factor,  the  reaction  constant,  or  velocity 
constant. 

50.  Let  us  assume  that  the  molecules  of  a  compound  are  dis- 
sociated by  heat  into  two  others,  the  process  being  expressed  by 


in  which  A,  B  and  C  represent  single  molecules.  Of  the  sub- 
stance A,  a  gram-molecules  were  originally  present,  but  in  the 
course  of  a  definite  time,  t,  x  gram-molecules  have  undergone  the 
above  decomposition.  The  problem  is  to  express  the  reaction 
velocity  at  any  moment.  At  the  beginning  (during  the  first 
minute)  the  reaction  velocity  is  proportional  to  a;  after  the  time 
t,  when  the  concentration  has  fallen  to  a  —  x,  the  amount  converted 
during  the  succeeding  minute  will  be  proportional  to  a—  x.  The 
reaction  velocity  thus  constantly  diminishes.  This  being  the  case,  it 
is  evident  that  s=k-a  and  s'  =k(a—  x)  are  not  the  true  expressions, 
respectively,  for  the  reaction  velocity  in  the  first  minute  and  in  the 
minute  following  the  time  t.  They  would  be  correct,  if  the  con- 
centration remained  constant  during  these  minutes  instead  of 


76  INORGANIC  CHEMISTRY,  [§§50- 

diminishing,  as  it  does  in  reality.  However,  we  can  approach 
the  real  velocity  by  considering  not  one  minute  but  a  very  small 
fraction  of  this  unit  of  time,  which  we  will  call  At;  the  smaller  At 
is  taken  the  less  is  the  concentration  change.  Supposing  that 
the  quantity  of  A  which  is  transformed  in  this  very  short  period 

Ax 
At  is  Ax,  the  expression  -j-  must  be  very  close  to  the  real  velocity, 

because  it  indicates  the  quantity  transformed  in  a  unit  of  time 
so  small  that  the  concentration  scarcely  diminishes  during  it. 
We  approach  the  true  velocity  nearer  and  nearer  according  as 
we  take  At  smaller  and  smaller)  and  when  it  is  made  infinitely 

Ax 
small  -T-  becomes   the   exact  expression  of  the  velocity.     It  is 

customary  to  express  such  infinitely  small  quantities  by  the  letter 
d,  thus:  —  .  The  mathematical  expression  for  the  velocity  at 
a  time  t,  when  the  concentration  is  a—  x,  thus  becomes 

*-£-*(«-,),     .. 

K  being  the  velocity  constant. 

The  above  reaction  is  termed  unimolecular.  When  two  (like  or 
different)  molecules  react  with  each  other,  the  reaction  is  called 
bimolecular;  it  may  be  represented  by  the  equation 

A+B=C    or     =C+Z>  +  ... 

The  equation  for  the  reaction  velocity  is  different  in  this  latter 
case.  Assuming  that  originally  a  gram-molecules  of  A  and  b  gram- 
molecules  of  B  take  part  in  the  reaction,  and  that  x  gram-molecules 
of  A  and  of  B  are  decomposed  at  the  end  of  the  period  t,  there  must 
be,  respectively,  a—x  and  b—  x  of  the  two  substances  present  at 
this  moment.  The  reaction  velocity  will  then  be  proportional  to 
the  product  of  these  quantities,  thus: 


in  which  Kf  is  again  a  constant  ;  for  suppose  that  there  were  only 
one  molecule  of  A  present;   the  possibility  of  its  reacting  with  a 


51.]  DISSOCIATION.  77 

molecule  of  B  would  then  be  proportional  to  the  number  of  mole- 
cules of  B.  When  there  are  a  —  x  molecules  of  A  this  possibility 
becomes  a—x  times  as  large. 

It  is  assumed  in  the  above  that  the  temperature  is  constant. 
We  shall  see  in  §  104  that  this  factor  has  a  great  influence  on  the 
reaction  velocity. 

51.  The  manner  of  expressing  the  condition  of  equilibrium  is 
now  plain.  Assuming  that  the  reaction  velocity  of  the  one  system 
is  S  and  that  of  the  other  S',  equilibrium  must  exist  when 

S=S'. 

The  state  of  equilibrium  may  therefore  be  defined  as  that  state 
in  which  the  reaction  velocities  of  both  systems  have  become  equal. 

Let  us  apply  these  considerations  to  the  dissociation  of  hydro- 
gen iodide.  This  may  be  expressed  by 


If  a  gram-molecules  HI  per  unit  volume  are  present  originally  and 
f  of  these  are  decomposed  after  a  given  period,  the  reaction  velocity 
at  this  moment  (since  in  this  case  a  =  b,  and  the  reaction  is  evi- 
dently bimolecular)  is 

S=C(a-x)2, 
C  being  a  constant. 

x 
From  the  x  gram-molecules  of  hydrogen  iodide  —  gram-mole- 

cules of  hydrogen  and  an  equal  amount  of  iodine  have  been  formed. 
The  velocity  of  formation  of  HI  from  H2  and  I2  is  therefore  ex- 
pressed by  the  equation 


in  which  C'  is  a  constant.     Accordingly  equilibrium  will  exist 
when 


4C 
in  which  K  is  substituted  for  -. 


78  INORGANIC  CHEMISTRY.  [§§  51- 

This  equation  may  be  written  in  a  slightly  different  way,  since 
in  gases  the  number  of  molecules  per  unit  volume  is  proportional 
to  the  pressure  (§  31).  Assuming  that  at  a  given  moment  the 
pressure  of  the  hydrogen  iodide  still  present  is  p,  that  of  the  hydro- 
gen is  plf  and  that  of  the  iodine  vapor  p2,  the  equilibrium  constant, 
K,  can  be  represented  by  the  expression 


It  has  been  ascertained  that  the  dissociation  is  less  if  the  hy- 
drogen iodide  was  originally  mixed  with  hydrogen  or  with  iodine 
vapor.  The  necessity  of  this  being  true  follows  immediately  from 
the  above  equation,  for  the  addition  of  these  gases  amounts  to  an 
increase  of  pi  or  p2.  UK  is  to  remain  constant,  p,  or  in  other 
words  the  mass  of  undissociated  hydrogen  iodide,  must  increase. 
We  see  also  from  the  equation  that  the  same  increase  of  H2  or  of 
12  must  have  the  same  influence  on  the  equilibrium.  A  further 
conclusion  from  this  equation  is  that,  when  p,  pi  and  p2  are  in- 
creased n-fold,  i.e.  when  hydrogen  iodide  undergoing  dissociation 
is  compressed  or  expanded  at  a  constant  temperature,  the  degree 
of  dissociation  must  remain  unaltered,  since 


(np)2          p2 

This,  too,  is  confirmed  by  experiment. 

In  the  dissociation  of  hydrogen  iodide  the  gas  volume  does  not 
change,  since  two  molecules  (2HI)  yield  two  molecules  (H2  and  I2). 
In  all  such  cases  the  degree  of  dissociation  must  be  independent 
of  the  volume,  because  an  increase  or  decrease  in  the  latter  causes 
changes  in  the  concentration  of  the  reacting  gases  which  are  pro- 
portional to  each  other,  and  hence  the  factor  representing  the 
concentration  falls  out  of  the  equation. 

Hydrogen  iodide  is  also  decomposed  by  light.  It  is  a  peculiar  fact 
that  this  dissociation  is  unimolecular  (HI  =  H  +  1)  ,  while  that  caused  by 
rise  of  temperature  is  bimolecular.  This  may  be  demonstrated  by  the 
following  very  general  method.  When  the  reaction  is  unimolecular, 

the  equation  for  the  velocity  of  decomposition  is  ^=K(a—x).    When, 


52.]  FLUORINE.  79 


however,    it    is    bimolecular    (2HI  =  H2  +  I2),    the    equation    becomes 
jT=K(a—  z)2.     With  the  help  of  integra 
be  solved  for  K'}  from  the  first  we  find 


T=K(a—  z)2.     With  the  help  of  integral  calculus  these  equations  can 


where  loge  is  the  natural  logarithm,  and  from  the  second  : 


a(a-x) 


If  now  we  determine  x  for  various  values  of  t,  the  values  of  K  can  be 
calculated;  they  must  be  constant.  If  this  constancy  appears  in  (1), 
the  reaction  is  unimolecular,  if  in  (2),  bimolecular. 


FLUORINE. 

52.  This  element  was  first  isolated  from  its  compounds  by 
MOISSAN  in  1886.  It  occurs  in  nature  chiefly  in  combination  with 
calcium  as  fluor  spar,  CaF2,  and  in  certain  rare  minerals. 

The  great  difficulty  in  obtaining  it  in  the  free  state  is  due  to 
its  very  great  affinity,  which  makes  it  unite  with  other  elements 
even  at  ordinary  temperatures.  As  yet  it  has  only  been  success- 
fully prepared  by  the  electrolysis  of  pure  anhydrous  hydrofluoric 
acid  in  which  potassium  fluoride  has  been  dissolved  to  make  the 
liquid  a  conductor. 

The  manner  in  which  MOISSAN  accomplished  this  is  interesting. 
A  mixture  of  about  200  g.  anhydrous  hydrofluoric  acid  and  60  g. 
hydrogen  potassium  fluoride  is  introduced  into  a  copper  U-tube  (Fig. 
22)  of  a  capacity  of  about  300  c.c.,  which  has  two  lateral  exit-tubes. 
The  open  ends  of  the  U-tube  are  closed  with  stoppers  FF,  made  of 
fluor  spar  and  wrapped  in  very  thin  sheet  platinum. 

The  cylindrical  electrodes  it  of  platinum-iridium  pass  through  the 
stoppers  and  are  held  in  place  by  the  copper  screws  EE,  which  fit  tightly 
to  the  ends  of  the  U-tube,  with  the  help  of  a  band  of  lead  P. 

During  the  electrolysis  the  apparatus  is  kept  at  the  constant  tem- 
perature of  —23°  (by  boiling  methyl  chloride).  The  free  fluorine, 
which  is  given  off  as  a  gas  at  the  positive  electrode,  is  first  passed  through 
a  platinum  vessel  that  is  cooled  by  a  mixture  of  solid  carbon  dioxide 
and  alcohol,  in  order  to  condense  the  acid  fumes  which  were  carried 


80 


INORGANIC  CHEMISTRY. 


[§§  52- 


over  with  it.  The  last  traces  of  the  acid  are  removed  by  conducting 
the  gas  through  two  platinum  tubes  containing  sodium  fluoride,  which 
absorbs  the  hydrofluoric  acid.  The  free  fluorine  gas  was  collected  by 
MOISSAN  in  a  platinum  tube,  whose  two  ends  were  closed  with  plates 
of  fluor  spar  so  that  one  could  look  through. 


FIG.  22. — PREPARATION  OF  FLUORINE  BY  ELECTROLYSIS.     (AFTER  MOISSAN.) 

Later  MOISSAN  found  that  perfectly  pure  fluorine  attacks  glass  but 
very  slowly,  so  that  the  gas  may  be  collected  in  glass  vessels. 

Physical  Properties. — Fluorine  is  a  gas  with  a  very  pungent 
odor  and  a  greenish-yellow  color,  which  is  somewhat  paler  than 
that  of  chlorine.  As  a  liquid  it  boils  at  — 187°  and  is  bright  yellow. 
•It  can  be  condensed  in  a  glass  vessel.  When  cooled  by  liquid 
^hydrogen  it  freezes  to  a  white  mass,  that  melts  at  —223°.  The 
specific  gravity  of  the  gas  is  19  (O=16),  that  of  the  liquid  1.14 
(water  =1). 

Chemical  Properties. — Of  all  the  elements  now  known  fluorine 
has  the  strongest  tendency  to  form  compounds.  It  combines 
with  hydrogen  in  the  dark  at  ordinary  temperatures  in  an  explo- 
sive manner.  MOISSAN  demonstrated  this  with  the  help  of  the 
above  apparatus  by  reversing  the  electric  current  while  fluorine 


53.]  FLUORINE.  81 

was  being  generated;  thus  a  mixture  of  hydrogen  and  fluorine  was 
formed,  which  at  once  exploded.  As  low  as  —252.5°,  solid  fluorine 
unites  with  liquid  hydrogen  immediately,  producing  a  flame. 
Finely  divided  carbon  ignites  instantaneously  in  fluorine  gas,  form- 
ing CF4.  With  sulphur,  red  phosphorus,  lime  and  other  sub- 
stances fluorine  reacts  vigorously  even  at  — 187°.  Fluorine  com- 
bines with  most  metals  instantly  and  violently;  it  does  not 
unite  with  oxygen,  even  when  it  is  heated  with  the 
latter  to  500°  or  mixed  in  the  liquid  state  with  liquid  oxygen  at 
—  190°.  The  alkali  metals  (potassium  and  sodium)  and  the 
alkaline-earth  metals  (calcium,  strontium,  barium)  take  fire  in 
fluorine  gas  at  ordinary  temperatures  with  the  formation  of 
fluorides.  Finely  divided  iron  glows  faintly  in  it.  Copper  be- 
comes covered  with  a  layer  of  copper  fluoride,  CuF2,  which  pro- 
tects it  against  farther  corrosion ;  hence  the  possibility  of  employ- 
ing this  metal  for  fluorine  generators.  Gold  and  platinum  are  not 
attacked  by  fluorine, — a  rather  striking  fact,  since  these  metals 
are  acted  on  by  chlorine,  which  otherwise  displays  a  weaker 
chemical  affinity. 

Fluorine  reacts  readily  with  hydrogen  compounds;  e.g.  water 
is  decomposed  by  it  at  ordinary  temperatures  into  hydrofluoric 
acid  and  strongly  ozonized  (as  high  as  14%  by  volume)  oxygen. 
It  sets  chlorine  free  from  potassium  chloride,  forming  potassium 
fluoride. 

The  molecule  of  gaseous  fluorine  is  expressed  by  the  formula  F2. 
Its  vapor  density  being  19,  the  molecular  weight  is  38.  Inasmuch 
as  no  fluorine  compound  contains  less  than  19  g.  fluorine  per 
gram-molecule,  but  frequently  a  multiple  of  this  amount,  the 
atomic  weight  of  fluorine  becomes  19  and  its  molecular  formula  Y& 

HYDROGEN  FLUORIDE,  or  HYDROFLUORIC  ACID,  HF. 

53.  This  compound  was  discovered  by  SCHEELE  in  1771  upon 
heating  together  fluor  spar  and  sulphuric  acid : 

CaF2  +  H2SO4 = CaS04 + 2HF. 

This  is  still  the  usual  method  of  preparing  the  substance.  A 
mixture  of  powdered  fluor  spar  and  dilute  sulphuric  acid  is  distilled 
in  an  apparatus  of  platinum  or  lead,  since  glass  is  instantly 
attacked  by  hydrofluoric  acid.  The  distillate  is  an  aqueous  solu- 


82  INORGANIC  CHEMISTRY.  [§§53- 

tion  of  the  acid,  which  for  the  above  reason  must  be  preserved  in 
bottles  of  lead  or  caoutchouc. 

By  direct  synthesis  from  its  elements  (§  52)  hydrofluoric  acid 
may  also  be  obtained.  Another  method  is  by  the  action  of  hydro- 
gen on  a  fluorine  compound;  e.g.  silver  fluoride,  when  heated 
in  a  current  of  hydrogen,  gives  hydrogen  fluoride. 

Still  other  methods  are  "by  the  action  of  fluorine  on  hydrogen 
compounds  (§  52)  and  by  the  direct  decomposition  of  certain  com- 
pounds, such  as  hydrogen  potassium,  fluoride,  KF-HF,  which 
splits  up  on  heating  into  the  two  fluorides.  This  last  reaction  is 
made  use  of  when  anhydrous  acid  is  sought. 

Physical  Properties. — Anhydrous  hydrofluoric  acid  is  a  color- 
less liquid  at  ordinary  temperatures.  It  boils  at  19.5°  and  solidifies 
at  - 102.5°.  Sp.  g.  (H=  1)  =  0.9879  at  15°.  It  has  an  extremely 
pungent  odor  and. is  very  poisonous  when  inhaled.  It  is  very 
soluble  in  water. 

Chemical  Properties.  —  The  aqueous  solution  of  hydrogen 
fluoride,  the  " hydrofluoric  acid"  of  commerce,  possesses  entirely 
the  character  of  an  acid;  it  evolves  hydrogen  with  most  metals, 
the  precious  metals,  however,  and  also  lead,  being  unaffected  by  it. 
The  fluorides  of  the  metals  are,  in  general,  soluble  in  water;  some, 
however,  such  as  those  of  copper  and  lead,  dissolve  with  difficulty, 
while  those  of  the  alkaline  earths  (Ca,  Sr  and  Ba)  are  insoluble. 
It  is  a  peculiar  characteristic  of  the  alkali  fluorides  that  they  are 
able  to  combine  with  a  molecule  of  the  acid,  forming  double  fluorides 
like  that  described  above,  KF-HF. 

This  characteristic  is  probably  due  to  the  fact  that  in  aqueous  solu- 
tion the  molecule  of  hydrofluoric  acid  is  IT2F2.  The  formation  of  such 
double  molecules  is  often  observed  for  acids  (especially  organic  acids) .  It 
is  called  association.  Thus  liquid  water  consists  in  all  probability  of 
H4O2  molecules. 

The  most  important  property  of  the  gas  for  practical  purposes 
is  that  it  attacks  glass  (cf.  §  193).  As  a  result  it  finds  extensive 
use  in  etching  glass. 

Glass  may  be  etched  in  two  ways — with  a  solution  of  the  gas  or  with 
the  gas  itself.  In  the  first  case  the  etching  is  shiny  and  transparent; 
in  the  second  dull.  The  glass  object  is  covered  with  a  coat  of  wax  in 
which  the  figures  or  letters  which  one  desires  etched  on  the  glass  may 
be  drawn  with  a  stylus.  Then  the  object  is  either  dipped  in  dilute 


54.]  HYDROFLUORIC   ACID.  33 

hydrofluoric  acia  for  a  while  or  set  over  a  leaden  dish  which  contains  a 
mixture  of  sulphuric  acid  and  calcium  fluoride  kept  slightly  warm  by  a 
low  flame.  Only  the  places  where  the  coating  was  removed  are  attacked, 
so  that,  when  the  latter  is  subsequently  dissolved  off  (by  turpentine  or 
alcohol),  the  etch-figure  is  visible. 

MOISSAN  has  proved  that  glass  is  also  attacked  by  perfectly  drv 
hydrofluoric  acid  gas. 

The  formula  of  hydrofluoric  acid  gas  is  HF,  which  can  be  deter- 
mined in  exactly  the  same  way  as  was  done  for  the  analogous 
chlorine  and  bromine  compounds. 

Compounds  of  the  Halogens  with  each  other. 

54.  The  halogens,  or  salt-formers,  i.e.  the  elements  fluorine,  chlorine, 
bromine  and  iodine  (so-called  because  they  form  salts  with  metals  by  direct 
combination),  can  unite  with  each  other  to  form  rather  unstable  compounds. 
In  general  the  most  stable  of  these  compounds  are  those  whose  component 
halogens  show  the  greatest  dissimilarity. 

Iodine  unites  with  fluorine  to  form  a  compound  IF5,  which  can  exist  even 
in  the  gaseous  state.  There  is  also  a  BrF3 ;  chlorine  and  fluorine,  however, 
do  not  combine  with  each  other. 

Chlorine  and  bromine  at  low  temperatures  give  an  unbroken  series  of 
mixed  crystals  (§  212,  2),  but  form  no  compound.  With  iodine  chlorine  gives 
two  compounds,  IC1  and  IC13.  It  depends  on  the  quantity  of  chlorine  present 
as  to  whether  the  former  or  the  latter  is  obtained.  IC1  is  a  reddish  brown  oil 
that  eventually  yields  crystals  melting  at  24.7°;  it  boils  at  101.3°.  Water 
decomposes  it  into  iodic  acid,  iodine  and  hydrogen  chloride.  It  exists  in  two 
modifications.  IC13  crystallizes  in  long  yellow  needles  and  on  fusing  dis- 
sociates almost  completely  into  IC1  and  C12.  In  a  small  quantity  of  water 
it  dissolves  almost  unchanged ;  but  a  larger  quantity  of  water  decomposes  it 
partially  into  hydrogen  chloride  and  iodic  acid. 

Bromine  and  iodine  give  only  one  compound,  BrI,  which  is  considerably 
dissociated  in  the  liquid  as  well  as  in  the  gaseous  state. 

Oxygen  Compounds  of  the  Halogens. 

With  the  exception  of  fluorine,  the  halogens  are  known  to  form 
various  oxygen  compounds,  having  the  common  property,  of 
instability,  i.e.  of  being  easily  decomposed.  Most  of  them  can 
combine  with  water,  forming  acids.  Oxides  which  show  this  latter 
property  are  called  acid  anhydrides.  The  acids  which  are  thus 
formed  from  the  halogen  oxides  contain  each  but  one  hydro- 
gen atom,  and  this  can  be  replaced  by  a  metal.  Acids  contain- 
ing one  hydrogen  atom  which  can  be  thus  substituted  are  called 
monobasic. 


84  INORGANIC  CHEMISTRY.  [§§  55- 

HYPOCHLOROUS  OXIDE.      CHLORINE  MONOXIDE,  C120. 

55.  This  compound  can  be  prepared  by  passing  chlorine  bver 
dry  mercuric  oxide  at  a  low  temperature: 

2HgO  +  2C12  =  C12O  +  HgO  •  HgCl2. 

Hypochlorous  oxide  is  a  brownish-yellow  gas  at  ordinary  tem- 
peratures. It  can  be  condensed  by  strong  cooling  to  a  dark-brown 
liquid,  which  boils  at  +  5°.  It  is  an  extremely  dangerous  sub- 
stance, especially  in  the  liquid  state,  since  the  slightest  mechanical 
disturbances  make  it  explode  vigorously,  breaking  up  into  its 
elements.  It  is  possible  to  distil  it  without  decomposition,  only 
when  everything  with  which  it  comes  in  contact  is  entirely  free 
from  dust  (organic  matter).  It  acts  upon  sulphur,  phosphorus 
and  compounds  of  carbon  with  explosive  violence. 

The  composition  of  this  compound  was  determined  by  BALARD 
in  the  following  way:  He  introduced  50  vols.  of  the  gas  into  a 
tube  over  mercury  and  decomposed  it  by  gently  warming.  He 
thus  obtained  a  mixture  of  chlorine  and  oxygen  which  occupied 
somewhat  less  than  75  vols.  After  the  chlorine  was  removed  by 
caustic  potash,  25  vols.  remained,  i.e.  50  vols.  chlorine  were  present, 
the  slight  difference  which  was  observed  being  ascribable  to  the 
fact  that  a  little  chlorine  had  united  with  the  mercury  in  the  tube, 
1  voL  hypochlorous  oxide  yielded  therefore  1  vol.  chlorine  and 
J  voL  oxygen.  This  indicates  the  formula 


2C12O=2C12 

2  vols.        2  vols.      1  vol. 

The  vapor  density  of  the  compound  was  found  to  be  3.03 
(air  =  1),  or  43.63  (O  =  16).  Its  molecular  weight  therefore  becomes 
87.26,  corresponding  to  the  formula  C12O  (201  =  71;  O  =  16;  sum  = 
87). 

HYPOCHLOROUS  ACID,  HC10. 

56.  When  chlorine  monoxide,  C120,  is  passed  into  water,  it  is 
absorbed;  the  solution  contains  hypochlorous  acid: 

=2HC1O. 


56.]  HYPOCHLOROUS  ACID.  85 

This  compound  is  known  only  in  aqueous  solution.  Its  com- 
position is  studied  in  its  salts! 

The  same  aqueous  solution  can  also  be  obtained  by  adding 
finely  powdered  mercuric  oxide  to  chlorine  water. 

HgO  +  2C1 2  +  H  20  =  HgCl  2  +  2C10H. 

Soluble. 

Upon  distillation  a  pure  aqueous  solution  of  the  acid  is 
obtained. 

Still  another  method  of  preparing  the  acid  solution  is  to  lead 
chlorine  into  the  solution  of  a  base,  e.g.  potassium  hydroxide,  at 
the  ordinary  temperature,  whereupon  a  salt  of  hypochlorous  acid 
(hypochlorite)  is  formed: 

2KOH  +C12  =  KC1 +KC10  +H2O. 

By  carefully  treating  the  hypochlorite  with  the  equivalent 
amount  of  nitric  acid  the  hypochlorous  acid  is  set  free  and  can  be 
separated  from  the  salts  by  distillation. 

When  concentrated,  the  aqueous  solution  of  hypochlorous  acid 
has  a  golden  color.  It  is  unstable;  only  dilute  solutions  can  be 
distilled  without  decomposition.  It  oxidizes  vigorously,  breaking 
up  into  oxygen  and  hydrochloric  acid: 

2C1OH  =  2HC1  +  O2. 

On  the  addition  of  hydrochloric  acid  all  the  chlorine  of  both 
compounds  is  set  free: 

HC1O  +  HC1=C12  +  H2O. 

The  hypochlorites  act  just  like  the  free  acid,  since  the  presence 
of  very  weak  acids,  e.g.  the  carbonic  acid  of  the  air,  serves 
to  liberate  hypochlorous  acid.  They  are  therefore  extensively 
employed  as  bleaching  agents  (§  27).  A  solution  of  potassium 
hypochlorite  (eau  de  Javelle)  is  used  for  this  purpose,  but  chloride 
of  lime  ("  bleaching  powder/7  §  258)  deserves  particular  notice.  The 
latter  is  obtained  by  treating  lime  with  chlorine  at  ordinary  tern- 


86  INORGANIC  CHEMISTRY.  [§§  56- 

peratures.  The  bleaching  action  of  hypochlorous  acid  is  twice 
as  great  as  that  of  the  chlorine  which  it  contains  would  be,  if  the 
latter  were  to  act  in  the  free  state  : 

2C1+H2O  =  2HC1+O    and    2C10H  =  2HC1  +  20. 

However,  it  should  be  remembered  that  two  atoms  of  chlorine 
were  necessary  to  form  the  one  HC1O  molecule: 

2KOH  +  C12  =  KC1  +  KC1O  +  H2O. 

On  shaking  an  aqueous  solution  of  hypochlorous  acid  with  mercury 
a  brownish-yellow  precipitate  of  mercuric  oxychloride,  nHgO  •  HgCl2,  is 
formed,  which  is  insoluble  in  hydrochloric  acid.  Chlorine  water,  on  the 
other  hand,  when  shaken  with  mercury,  gives  white  mercuric  chloride, 
HgCl2  (sublimate).  These  reactions  enable  us  to  distinguish  between 
the  two  substances. 

In  a  dilute  aqueous  solution  of  chlorine  we  have  the  following  equilib- 
rium: 


as  is  shown  by  the  facts  that  the  solution  reacts  distinctly  acid  toward 
litmus  and  that  the  hypochlorous  acid  can  be  separated  from  the  hydro- 
chloric acid  by  distillation. 

The  difference  in  the  action  of  chlorine  water  and  a  solution  of  hypo- 
chlorous  acid  on  mercury  is  due  to  the  fact  that  in  the  above  equilibrium 
the  system  C12  +  H20  is  by  far  the  predominant  one. 

CHLORINE  DIOXIDE,  C102. 

57.  This  gas  is  formed  when  potassium  chlorate,  KClOa,  is 
treated  with  concentrated  sulphuric  acid.  Chloric  acid  is  at  first 
set  free  and  this  decomposes  as  follows: 

3HClp3  =  HC104  +  2C102  +  H20. 

Chloric          Perchloric 
acid.  acid. 

Chlorine  dioxide  is  a  dark-yellow  gas.  It  can  be  condensed  to 
a  liquid,  which  boils  at  9.9°  and  solidifies  at  —79°  to  a  yellow 
crystalline  mass.  It  has  a  peculiar  otlor  resembling  chlorine  and 
burned  sugar. 

Chlorine  dioxide  is  extremely  explosive;  warming,  jarring  or 
contact  with  organic  substances  causes  it  to  explode  with  vio- 
lence. Light  slowly  decomposes  it. 


57.]  CHLORINE  DIOXIDE.  37 

The  following  experiments  give  one  an  idea  of  the  vigor  with  which  it 
causes  oxidation.  (1)  When  finely  powdered  sugar  is  mixed  carefully 
with  potassium  chlorate  and  a  drop  of  concentrated  sulphuric  acid  is 
added,  the  whole  mass  bursts  into  flame.  The  chlorine  dioxide  set  free 
makes  the  sugar  burn  at  ordinary  temperature.  (2)  Place  a  few  pieces 
of  yellow  phosphorus  and  some  crystals  of  potassium  chlorate  under 
water  and  allow  a  few  drops  of  concentrated  sulphuric  acid  to  flow  down  on 
the  two  substances.  The  phosphorus  at  once  burns  under  water  with 
a  brilliant  light. 

Chlorine  dioxide  is  soluble  in  water.  Such  a  solution  can  be 
easily  prepared  by  floating  a  little  porcelain  cup  in  a  large  crystal- 
lizing-dish  with  a  flat  brim  and  containing  220  c.c.  water,  putting 
into  the  cup  12  g.  potassium  chlorate  and  adding  a  cooled  mixture 
of  44  c.c.  concentrated  sulphuric  acid  and  11  c.c.  water.  The 
crystallizing-dish  is  then  covered  with  a  glass  plate.  The  chlorine 
dioxide  evolved  dissolves  in  the  water,  forming  a  yellow  solution. 

When  a  base  is  added  to  a  chlorine  dioxide  solution,  a  chlorite 
(§  58)  and  a  chlorate  are  formed: 

2KOH     +     2C1O2     =     KC1O2     +     KC103     +     H20. 

Pot.  chlorite.          Pot.  chlorate. 

This  reaction  proceeds  very  slowly  m  dilute  solution. 

The  composition  of  chlorine  dioxide  was  determined  by  GAY- 
LUSSAC  as  follows:  He  allowed  the  gas  to  flow  through  a  capillary 
tube  with  three  bulbs.  By  heating  the  part  of  the  tube  in  front 
of  the  bulbs  he  decomposed  the  gas,  the  action  being  non-explo- 
sive in  so  narrow  a  space.  Thus  there  was  obtained  in  the  bulbs 
a  mixture  of  oxygen  and  chlorine  in  the  same  proportions  as  they 
are  contained  in  the  compound.  The  chlorine  was  absorbed  by 
potash  and  the  residual  gas  (oxygen)  was  passed  over  into  a 
measuring-tube.  The  capacity  of  the  bulbs  being  known,  it  was 
possible  from  these  data  to  calculate  the  volume  ratio  of  oxygen 
and  chlorine.  It  was  found  that  2.00  vols.  of  the  oxide  yield  .987 
vol.  chlorine  and  2.063  vols.  oxygen.  The  combining  ratio  is  very 
close  to  that  of  1 : 2,  represented  by  the  formula  C1O2 : 
2C1O2  =  C12  +  202. 

2  vols.        1  vol.       2  vols. 

This  formula  is  also  confirmed  by  the  vapor  density,  which  was 
found  to  be  34.5  at  10.5°,  while  the  formula  C102  demands 
35.5+2X16-  _ 


88  INORGANIC  CHEMISTRY.  [§§58- 


CHLOROUS  ACID,  HC10,. 

58.  This  acid  is  unknown  in  the  pure  state.     Its  sodium  salt  is  formed 
by  the  action  of  sodium  peroxide  solution  on  a  chlorine  dioxide  solution: 

2C102  +  Na202  =  2NaC102  +  O2. 

The  silver  salt,  AgC102,  is  a  yellow  crystalline  powder,  as  is  also  the  lead 
salt,  Pb(C102)2;  they  are  both  difficultly  soluble  in  water,  and  break  up 
even  on  warming  to  100°  in  an  explosive  manner.  The  anhydride  of 
chlorous  acid,  corresponding  to  the  formula  C12O3,  is  not  known. 

CHLORIC  ACID,  HC103. 

59.  The  chlorates  of  potassium  or  barium  are  the  usual  start- 
ing-points   for    the    preparation    of    chloric    acid.     When    dilute 
sulphuric  acid  is  added  to  the  solution  of  the  barium  chlorate, 
barium  sulphate  is  precipitated  and  a  dilute  solution  of  chloric 
acid  is  obtained,  which  may  be  filtered  off  from  the  sulphate  and 
dried  in  a  vacuum  desiccator  over  concentrated  sulphuric  acid. 
In  this  way  a  40%  solution  of  the  acid  may  be  obtained.     On 
concentrating  it  any  farther,  decomposition  takes  place,  oxygen 
being    evolved    and    perchloric    acid    formed.     The    concentrated 
acid  is  a  powerful  oxidizing  agent;   wood  or  paper  ignites  when 
brought  in  contact  with  it.     It  oxidizes  hydrochloric  acid,  chlo- 
rine being  given  off;    further  sulphuretted  hydrogen,  sulphurous 
acid  and  others,  even  in  dilute  solution.     The  following  reaction 
is  very  characteristic  of  chloric  acid.     When  indigo  solution  is 
added  to  a  dilute  solution  of  the  acid,  the  former  is  not  decolorized ; 
however,  on  the  addition  of  a  little  sulphurous  acid  the  color  dis- 
appears, since  the  chloric  acid  is  thereby  reduced  to  lower  oxides. 

The  salts  are  all  soluble  in  water,  that  of  potassium  being 
somewhat  difficultly  so,  however. 

The  composition  of  chloric  acid  was  ascertained  by  STAS  from 
an  analysis  of  silver  chlorate.  An  accurately  weighed  amount  of 
the  latter  was  reduced  by  a  solution  of  sulphurous  acid  to  silver 
chloride  and  this  was  filtered  off  and  weighed.  Since  he  knew 
from  previous  investigations  the  exact  composition  of  silver 
chloride,  the  analysis  of  the  silver  chlorate  was  complete. 
STAS  found  thus  that  silver  chlorate  consists  of 


60.]  PERCHLORIC  ACID.  89 

Silver 56.3948% 

Chlorine 18.5257% 

Oxygen 25.0795% 


Total  ........................  100.0000% 

The  atomic  weight  of  silver  is  107.88;  that  of  chlorine  35.46; 
that  of  oxygen  16.00.  We  then  find  that  the  ratio  of  the  atoms 
in  this  salt  is 

56.3948  18.5257  25.0795 

107\88~~  ~^46~-        •'  ~l6M~~ 

i.e.  very  close  to  1:1:3,  from  which  it  follows  that  the  empirical 
formula  of  the  salt  is  AgClOa,  that  of  the  acid  itself 


PERCHLORIC  ACID,  HC104. 

60.  This  compound  is  obtained  by  distilling  potassium  per- 
chlorate  with  an  excess  of  sulphuric  acid  of  96-97.5%  in  vacuo: 

KC104  +  H2S04  =  KHS04  +  HC104. 

Pure  perchloric  acid  boils  at  39°  under  a  pressure  of  56  mm. 
Hg,  and  has  a  specific  gravity  of  D422  =  1.764  at  22°.  It  is  a 
colorless  liquid  which  does  not  solidify  on  being  cooled  with  solid 
carbon  dioxide  and  alcohol  (about  —80°).  It  decomposes  slowly, 
taking  on  a  dark  color.  With  water  it  forms  different  hydrates; 
the  best  known  of  them  is  the  monohydrate,  HC1O4-H20,  which 
melts  at  50°;  with  more  water  a  thick  oily  liquid  is  formed, 
similar  to  concentrated  sulphuric  acid.  Such  an  acid  contains 
71.6%  HC1O4;  it  distils  without  change  in  composition  at  203° 
and  has  a  specific  gravity  of  1.82.  The  dilute  solution  of  the 
perchloric  acid  is  stable. 

In  the  concentrated  state  perchloric  acid  is  a  very  strong 
oxidizing  agent.  When  a  little  is  dropped  on  wood  or  paper, 
these  ignite  with  explosion.  Very  painful  flesh-wounds  are  pro- 
duced by  it.  When  dilute,  it  does  not,  however,  release  its  oxy- 
gen nearly  so  readily  as  chloric  acid.  It  can,  for  example,  be 
gently  warmed  with  hydrochloric  acid  without  giving  off  chlorine, 
and  it  is  not  reduced  by  sulphurous  acid.  By  these  facts  and 
by  its  yielding  no  chlorine  dioxide  with  sulphuric  acid  it  may 
be  distinguished  from  chloric  acid. 


90  INORGANIC  CHEMISTRY.  [§§  60- 

The  salts  of  perchloric  acid,  perchlorates,  are  all  solu- 
ble in  water;  that  of  potassium  and  especially  that  of  rubidium 
are,  however,  very  difficultly  soluble  in  cold  water. 

The  composition  of  perchloric  acid  has  been  determined,  as  in 
the  case  of  chloric  acid,  by  the  analysis  of  a  salt,  in  this  instance 
the  potassium  salt.  A  weighed  amount  of  the  latter  is  heated  to 
drive  off  all  the  oxygen.  The  loss  in  weight  indicates  the  amount 
of  the  latter.  The  analysis  of  the  remaining  potassium  chloride, 
KC1,  shows  the  amounts  of  potassium  and  chlorine.  From  these 
data  it  is  found,  in  the  same  manner  as  with  chloric  acid,  that  the 
empirical  formula  of  the  salt  is  KC104,  that  of  the  acid,  therefore, 
HC104. 

Chlorine  heptoxide,  C12O7,  is  the  oxide  corresponding  to  perchloric 
acid: 

2HC104-H20=C1207. 

It  may  be  obtained  by  slowly  adding  perchloric  acid  to  phosphorus  pent- 
oxide  cooled  below  —  10°.  By  distillation  on  a  water  bath  the  oxide  is 
obtained  as  a  colorless  liquid,  which  boils  at  82°.  It  is  more  stable  than 
the  other  oxides  of  chlorine ;  it  neither  attacks  paper  nor  acts  on  sulphur 
or  phosphorus  in  the  cold. 

OXYGEN  COMPOUNDS  OF  BROMINE. 

61.  Although  no  compounds  with  oxygen  alone  are  known,  there  are 
two  oxygen  acids,  viz.,  hypobromous  and  bromic. 

Hypobromous  acid,  HBrO,  can  be  obtained  in  the  same  way  as  HC1O, 
namely,  by  shaking  up  bromine  water  and  mercuric  oxide  together.  The 
dilute  solution  can  be  distilled  in  vacuo,  and  has  properties  entirely  anal- 
ogous to  those  of  hypochlorous  acid. 

Bromic  acid,  HBrO3,  can  be  obtained  from  the  barium  salt  with  sul- 
phuric acid  or  from  the  silver  salt  with  bromine-water : 

5AgBr03  +  3Br2  +  3H20  =  5AgBr  +  6HBr03. 

Insol. 

It  is  also  formed  when  chlorine  is  passed  into  bromine-water: 
Br2  +  5C12  +  6H2O  -  2HBr03  +  10HC1. 

It  corresponds  in  its  behavior  with  chloric  acid.  Many  reducing-agents, 
such  as  hydrogen  sulphide  and  sulphurous  acid,  are  able  to  extract  all 
its  oxygen.  Most  of  its  salts  are  difficultly  soluble  in  water.  When 
heated,  they  give  up  all  their  oxygen. 


62.]  OXYGEN  COMPOUNDS   OF  IODINE.  91 

OXYGEN  COMPOUNDS  OF  IODINE. 

62.  When  iodine  is  introduced  into  a  cold  dilute  solution  of  caustic 
potash  or  soda,  a  colorless  liquid  is  obtained,  which  has  other  properties 
when  fresh  than  it  has  later.  When  freshly  prepared  it  decolorizes  indigo 
solution  and  iodine  is  liberated  on  the  addition  of  very  weak  acids. 
*Later  on  these  two  properties  disappear.  It  is  therefore  to  be  supposed 
that  a  hypo-iodite  KIO  is  first  formed,  and  that  this  is  changed  slowly 
to  KI  and  KI03.  At  the  boiling-point  the  change  takes  place  almost 
instantly. 

Iodine  pentoxide,  I2O5,  is  the  anhydride  of  iodic  acid,  since  it 
can  be  obtained  by  heating  this  acid  to  170°, 

2HI03=H20  +  I205, 

and  yields  the  same  acid  when  dissolved  in  water.  It  is  a  white 
crystalline  substance,  which  breaks  up  into  its  elements  at  300°. 

Iodic  acid,  HIO3,  is  prepared  by  the  oxidation  of  iodine  with 
nitric  acid,  or,  better,  with  nitrogen  pentoxide. 

3I2  +  10HNO3  =  6HI03  +  10NO  +  2H20. 

Nitric  acid.  Nitric  oxide. 

Iodic  acid  is  crystalline  and  easily  soluble  in  water.  It  is  a  power- 
ful oxidizing-agent,  setting  free  chlorine  from  hydrochloric  acid, 
for  example. 

2HI03  +  10HC1  =  I2  +  5C12 + 6H20. 

It  reacts  instantaneously  with  hydriodic  acid,  all  the  iodine  of  both 
compounds  being  precipitated: 

5HI+HIO3=3H2O+6L 

The  salts  of  this  acid,  the  i  o  d  a  t  e  s  ,  are  in  general  not  very 
soluble  in  water;  however,  those  of  the  alkali  metals  dissolve 
rather  easily. 

On  heating  iodic  acid  with  concentrated  sulphuric  acid  oxygen  is 
evolved  and  the  compound  I204,  iodine  dioxide,  is  formed.     This  is  a 
lemon-yellow,  crystalline  powder  that  breaks  up  into  its  elements  above 
130°.     With  hot  water  it  reacts  quickly  to  form  iodine  and  iodic  acid: 
5I204+4H20=8HI03+I2. 

Periodic  acid,  HIC>4  +  2H2O,  is  formed  by  the  action  of  iodine 
on  perchloric  acid : 

HC1O4  + 1  +  2H20  =  HI04  •  2H2O  +  Cl. 

It  is  a  colorless  crystalline  solid  that  is  entirely  decomposed  at 
140°  into  iodine  pentoxide,  oxygen  and  water  (§  145). 


92  INORGANIC  CHEMISTRY.  [§§  63. 


NOMENCLATURE. 

63.  The  system  of  naming  the  various  halogen  oxygen-acids  is 
a  general  one,  which  is  also  used  for  the  acids  of  other  elements. 
The  best-known  acid  usually  has  the  suffix  -ic,  e.g.  chloric  acid,* 
phosphoric  acid,  sulphuric  acid,  etc.  Acids  that  contain  more 
oxygen  have  in  addition  the  prefix  per-,  thus  perchloric  acid 
and  persulphuric  acid.  Acids  containing  'less  oxygen  have 
the  suffix  -ous,  e.g.  chlorous  llcid,  sulphurous  acid,  phosphorous 
acid,  etc.  Those  which  contain  still  less  oxygen  have  the  suffix 
-ous  and  also  the  prefix  hypo-,  e.g.  hypochlorous  acid,  hyposul- 
phurous  acid  and  hypophosphorous  acid. 

The  names  in  use  in  pharmaceutical  chemistry  (see  the  National  Phar- 
macopoeia) follow  the  Latin.  Thus  we  have  Acidum  sulphuricum  (sul- 
phuric acid)  and  Acidum  sulphurosum  (sulphurous  acid). 

The  names  of  the  salts  of  the  best-known  (-ic)  acids  end  in 
-ate,  e.g.  potassium  chlorate,  -sulphate,  -phosphate.  The  salts  of 
the  -ous  acids  have  the  ending  -ite,  as  potassium  chlorite,  -sulphite, 
-phosphite.  The  salts  of  hypo-  -ous  acids  are  called  hypo-  -ites; 
thus  sodium  hypochlorite,  -hyposulphite,  -hypophosphite. 

The  names  of  the  anhydrides  correspond  to  those  of  their  acids.  < 
In  naming  oxides  the  name  of  the  element  with  or  without  the 
ending  -ic  is  used,  unless  there  is  more  than  one  oxide.  Where 
there  are  two  oxides,  the  name  of  the  one  with  the  more  oxygen 
ends  in  -ic,  that  of  the  other  in  -ous,  e.g.  mercuric  oxide,  arsenic 
oxide,  mercurous  oxide,  arsenious  oxide.  An  oxide  with  less 
oxygen  than  the  -ous  compound  is  given  the  prefix  hypo-,  and  one 
with  more  than  the  -ic  oxide  the  prefix  per-,  as  in  the  case  of  acids, 
thus  hypochlorous  oxide,  lead  peroxide.  In  some  cases,  for  the 
sake  of  euphony,  the  suffix  is  added  to  the  Latin  instead  of  the 
English  stem,  as  cuprous,  ferric,  etc. 

For  historical  reasons  many  names  now  in  use  do  not  conform  to  this 
system.  In  some  instances  the  oxide  first  discovered  took  the  suffix  -ic, 
and  those  subsequently  discovered  were  named  accordingly,  as  in  the 
case  of  the  nitrogen  oxides  (§119). 

It  is  not  uncommon  to  speak  of  oxides  of  the  general  formula  M2O3  as 
sesquioxides. 


64.J 


SUMMARY  OF  THE  HALOGEN  GROUP. 


93 


A  much  more  rational  system  is  to  indicate  the  number  of 
atoms  of  oxygen  by  the  Latin  or  Greek  numeral,  e.g.  chlorine 
protoxide,  or  monoxide,  iodine  pentoxide,  etc. 

SUMMARY  OF  THE  HALOGEN  GROUP. 

64.  It  is  evident  from  the  foregoing  descriptions  that  the 
properties  of  the  halogens  and  their  compounds  possess  great 
similarity  among  themselves.  A  closer  study  reveals  the  fact  that 
the  increase  of  atomic  weight  is  accompanied  by  a  gradual  change 
of  physical  and  chemical  properties.  For  example,  let  us  notice 
the  physical  properties: 


F. 

Cl. 

Br. 

I. 

19  0 

35  46 

79  92 

126  92 

Melting-point  

—  223° 

-102 

—  7 

+  113 

Boiling-point.  

—  187° 

-  33 

+  63 

+  200 

Sp.  g.  (liquid  or  solid).  .  . 
Color  

1  .  14(liquid) 
f  palegreen- 

1.33 
greenish- 

3.18 
brown. 

4.97 
violet- 

(  ish-yellow. 

yellow 

black 

It  is  seen  that  the  values  of  the  physical  constants  increase  on 
the  whole  with  the  atomic  weight.  The  purely  metalloid  char- 
acter of  the  first  three  is  also  found  in  iodine,  although  in  the  case 
of  the  latter  an  external  characteristic  of  metals,  viz.,  metallic 
lustre,  is  at  once  noticeable. 

The  affinity  for  hydrogen  decreases  as  the  atomic  weight 
increases.  We  saw  that  fluorine  combines  with  this  element  even 
in  the  dark  and  at  very  low  temperatures  in  an  explosive  manner; 
iodine  unites  with  hydrogen  directly  only  at  a  high  temperature 
and  the  compound  is  easily  decomposed  by  heat.  Inversely,  the 
oxygen  compounds  are  the  more  stable  the  higher  the  atomic 
weight  of  the  halogen.  While  a  halogen  of  low  atomic  weight 
displaces  a  halogen  of  higher  atomic  weight  from  its  hydroger 
compound,  e.g. 

HI  +  C1=HC1+I, 

the  halogen  with  higher  atomic  weight  can  on  the  other  hand 
replace  one  with  lower  in  its  oxygen  compounds,  setting  that 
other  one  free: 


94  INORGANIC  CHEMISTRY.  [§  65- 

ELECTROLYTIC  DISSOCIATION. 

65.  In  §  30  it  was  stated  that  the  properties  of  an  aqueous 
solution  of  hydrogen  chloride  differ  widely  from  those  of  the  dry 
gas.  It  was  also  stated  there  that  many  other  substances  undergo 
a  similar  change  of  properties  when  they  are  dissolved  in  water. 
We  may  now  consider  the  nature  of  this  change. 

If  we  investigate  the  freezing-point  depression  of  the  aqueous 
solution  of  an  acid,  base  or  salt  of  known  concentration,  we  find 
that  the  depression  does  not  correspond  with  that  calculated  from 
the  accepted  molecular  weight  (§  43).  The  freezing-point  depres- 
sion and  the  boiling-point  elevation  are  both  greater  than  they 
should  be.  A  1%  sodium  chloride  solution  would,  for  example, 

19 
be  expected  to  show  a  depression  of  -^-^=0.325°,  the  molecular 

Oo.O 

depression  for  water  (§  43)  being  19,  i.e.  AM  =19,  and  the  molec- 
ular weight  of  sodium  chloride  58.5.  In  reality,  however,  the 
depression  is  found  to  be  1.9  times  as  great,  namely,  0.617°.  As 
the  osmotic  pressure  is  proportional  to  the  freezing-point  depres- 
sion (§  42),  it  must  also  be  greater  than  the  calculated  amount. 

The  fact  at  once  occurs  to  us  that  gases,  to  which  dissolved 
substances  have  been  found  to  show  close  analogy,  also  exhibit 
a  similar  phenomenon.  In  numerous  instances  the  pressure 
exerted  by  a  definite  weight  of  gas  occupying  a  definite  volume 
at  a  definite  temperature  is  greater  than  the  calculation  indi- 
cates.— This  is  but  another  phase  of  the  observation  that  the 
vapor  density  of  some  gases  is  abnormally  low  at  certain  tem- 
peratures (§  47). — This  is  explained  by  assuming  a  breaking 
up  of  the  gas  molecules;  the  number  of  particles  free  to  move 
about  is  thus  increased  and  accordingly  the  pressure  becomes 
greater.  This  phenomenon  is  known  as  dissociation  (§  49). 

In  the  case  of  abnormal  osmotic  pressure  we  are  led  to  a 
similar  explanation  by  assuming  that  the  molecules  are  split  up 
into  several  independent  particles.  A  difficulty  arises,  however, 
when  we  try  to  conceive  the  nature  of  this  division.  In  solu- 
tions of  salts  in  water  it  would  be  possible  to  assume  a 
hydrolytic  separation,  i.e.  into  free  base  and  free  acid  (p.  104), 
which  would  necessarily  be  complete  in  dilute  solutions  of  the  salts 
of  strong  acids  and  bases,  inasmuch  as  the  osmotic  pressure  of  such 
solutions  in  concentrations  of  TV  normal  (§  93),  for  instance. 


65.]  ELECTROLYTIC  DISSOCIATION.  95 

amounts  to  double  the  calculated  pressure.  There  are,  however, 
serious  objections  to  such  a  hypothesis.  In  the  first  place,  it  has 
never  yet  been  possible  to  separate  such  a  solution  by  diffusion  into 
the  free  base  and  free  acid  which  it  is  supposed  to  contain.  A 
second  and  still  more  serious  objection  is  that  an  acid  or  base  in 
an  aqueous  solution  by  itself  exerts  an  osmotic  pressure  greater 
than  that  calculated.  Here,  however,  hydrolytic  dissociation  is 
impossible. 

The  question  as  to  the  real  nature  of  the  division  has  found  its 
answer  in  a  consideration  of  the  relation  which  exists  between  the 
abnormal  osmotic  pressure  and  the  transmission  of  the  electric 
current.  ARRHENIUS  observed  that  only  those  substances  which 
conduct  the  electric  current  in  aqueous  solution,  namely,  acids, 
bases  and  salts,  show  the  above-mentioned  abnormalities  in  osmotic 
pressure.  When  these  substances  are  dissolved  in  another  liquid 
than  water,  the  resulting  solution  is  a  non-conductor,  but  at  the 
same  time  its  osmotic  pressure  again  assumes  the  normal.  These 
facts  enable  us  to  perceive  the  connection  between  the  apparently 
disconnected  phenomena  of  abnormal  osmotic  pressure  and  elec- 
trolytic conduction. 

In  order  to  understand  this  relation  it  is  necessary  to  know 
the  usual  explanation  of  electrolytic  conduction.  Let  us  take 
hydrochloric  acid  as  an  example.  Perfectly  dry  hydrochloric  acid 
gas  is  a  non-conductor,  as  is  also  perfectly  pure  water.  However, 
when  the  gas  is  dissolved  in  water,  a  solution  is  obtained  which 
transmits  electricity  very  well.  Evidently  a  certain  reaction  must 
have  resulted  from  the  mixing  of  the  water  and  the  hydrogen 
chloride.  We  were  led  to  surmise  this  above  (§  29),  when  it  was 
found  that  this  gas  solution  does  not  obey  HENRY'S  law.  Since 
during  the  transmission  of  the  current  the  hydrogen  chloride  is 
broken  up  into  hydrogen  and  chlorine  while  the  water  remains 
unchanged,  it  must  be  assumed  that  the  hydrogen  chloride  molecules 
are  the  ones  which  have  undergone  a  change. 

The  phenomena  of  electrolytic  conduction  now  find  their  com- 
plete explanation  in  the  assumption  that  the*  change  which  the 
hydrochloric  acid  underwent  consisted  in  a  separation  of  its  mole- 
cules into  electrically  charged  atoms  (ions)  (§  267).  This  separa- 
tion may  have  been  complete  or  partial,  the  extent  depending  upon 
the  concentration  among  other  things.  When  a  current  passes 
through  the  solution,  the  negatively  charged  chlorine  ions  (the 


96 


INORGANIC  CHEMISTRY. 


65- 


anions)  are  drawn  toward  the  positive  electrode  (anode);  they 
become  electrically  neutral  on  contact  with  the  latter  and  escape 
from  the  liquid.  Similarly  the  positively  charged  hydrogen  ions 
(cations)  wander  toward  the  negative  electrode  (cathode). 
In  this  way  conduction  goes  on,  the  undivided  molecules  having 
no  part  in  it.  This  division  of  the  molecules  is  known  as  electro- 
lytic dissociation,  or  ionization. 

The  existence  of  free  ions  in  the  solution  of  an  electrolyte 
is  demonstrated  by  OSTWALD  in  the  following  manner.  The  tube 
abed,  Fig.  23,  is  nearly  filled  with  dilute  sulphuric  acid.  The 


& " 
m 


FIG.  23. 

narrowed  portion  be  is  about  40  cm.  long.  A  rod  of  amalgamated 
zinc  is  lowered  into  a  to  serve  as  the  positive  electrode,  while  a 
platinum  wire  is  fused  into  d  at  p  for  a  negative  electrode.  If 
connection  is  made  with  a  battery  of  ten  accumulators,  there  is 
an  immediate  evolution  of  hydrogen  at  p.  The  passage  of  the 
current  through  the  liquid  results  in  the  formation  of  zinc  sul- 
phate around  the  bar  in  a: 

Zn+H2SO4  =  ZnS04+H2. 

Now  if  this  hydrogen  has  to  pass  through  be  to  p,  it  must  cover 
the  40  cm.  in  a  very  brief  space  of  time.  However,  it  has  been 
shown  both  by  investigations  which  cannot  be  described  here 
and  by  calculus  that  this  migration  would  take  many  hours.  The 
hydrogen  appearing  at  p  as  soon  as  the  circuit  is  closed  cannot, 
therefore,  come  from  a;  the  most  natural  explanation  is  to  sup- 
pose that  there  are  already  free  ions  in  the  neighborhood  of  p 
and  that  they  are  discharged  by  the  current  and  given  off  from 
the  liquid  as  free  hydrogen. 

TOLMAN  has  shown  that,  when  a  long  tube  containing  a  solu- 
tion of  an  alkali  iodide  is  rotated  as  the  spoke  of  a  wheel  at  3000- 


65.]  ELECTROLYTIC   DISSOCIATION.  97 

5000  revolutions  per  minute,  the  outer  end  becomes  negative 
with  respect  to  the  inner  end.  The  solution  must  therefore 
contain  positive  and  negative  components  which  can  move 
independently  of  each  other.  The  iodine  ions,  being  much 
heavier  than  the  alkali  ions,  would  naturally  accumulate  at  the 
outer  end. 

In  order  that  this  hypothesis  of  dissociation  into  ions  may  also 
account  for  abnormal  osmotic  pressure,  it  must  be  assumed  that 
the  ions  are  independent  particles,  just  as  free  to  move  as  the 
molecules  are  supposed  to  be.  The  number  of  freely  moving 
particles  in  the  same  volume  is  thus  increased.  Hence,  whether 
the  amount  of  ionization  is  calculated  from  the  electrical  con- 
ductivity or  from  the  osmotic  pressure,  the  result  should  be  the 
same  according  to  the  above  hypothesis.  This  is  found  to  be 
the  case. 

Supposing  that  every  molecule  yields  n  ions  by  the  dissociation  and 
that  the  dissociated  portion  of  the  whole  number  of  molecules  is  r,  the 
number  of  freely  moving  particles  is 


The  osmotic  pressure  must  therefore  be  [l  +  ?-(n—  1)]  times  as  great  as 
in  the  case  of  undivided  molecules.  If  this  pressure  p  is  p0  in  the  latter 
case,  then 

wherefore 

rt  — r>- 

(1) 

From  the  electrical  conductivity  we  are  able  to  find  the  value  of  r  in 
the  following  way:  As  the  dilution  becomes  greater,  the  molecular  con- 
ductivity increases.  By  this  term  we  mean  the  specific  conductivity  of 
the  solution  multiplied  by  the  number  of  liters  in  which  a  gram-molecule 
of  the  substance  is  dissolved  As  the  dilution  is  gradually  increased,  the 
molecular  conductivity  approaches  a  definite  limit.  Since  the  conductivity 
is  only  due  to  the  d  i  s  s  o  c  i  a  t  e  d  molecules,  it  may  be  assumed  that, 
when  this  limit  is  reached,  all  the  molecules  are  broken  up  into  ions.  If 
the  molecular  conductivity  for  infinite  dilution  is  represented  by  A^  and 
that  for  the  dilution  v  (1  gram-molecule  in  v  liters)  by  XV}  it  is  evident  that 


The  following  table  shows  the  agreement  of  the  values  calculated 


98 


INORGANIC  CHEMISTRY. 


[§§65- 


by  the  two  methods.  The  values  opposite  r0  were  calculated  from  the 
observed  freezing-point  depressions  and  those  opposite  re  from  the  con- 
ductivities of  the  salt  solutions.  The  concentration  throughout  is  1  g.  per 
liter. 


KC1. 

NH4C1. 

KI. 

NaNOs. 

f 

0  82 

0  88 

0  90 

0  82 

T  .  . 

0  86 

0  84 

0.92 

0.82 

66.  Ionic  Equilibrium. — In  a  case  of  electrolytic  dissociation  we 
have  an  equilibrium  to  deal  with,  namely,  that  between  the  un- 
dissociated  molecules  on  the  one  hand  and  the  ions  on  the  other. 
In  the  case  of  a  monobasic  acid  this  equilibrium  may  be  repre- 
sented by 


where  A' 
(cation). 


is  the  acid  radical  (anion)  and  H*  the  hydrogen  ion 
For  a  base  we  have 

MOH<=±M'+OH'. 


We  may  apply  here  the  equilibrium  equation  deduced  in  §  49. 
Given  a  gram-molecules  of  AH  per  unit-volume,  of  which  x  are 
divided  into  two  ions  each,  then  the  equilibrium  is  represented  by 

a-x=Kx2. 

From  this  equation  it  necessarily  follows  that  the  dissociation  is 
diminished  by  the  introduction  into  the  solution  of  a  substance  with 
like  ions  (just  as  the  addition  of  hydrogen  or  iodine  reduces  the 
dissociation  of  hydrogen  iodide  gas,  §  50).  This  effect  (which  is 
called  the  "  common  ion  effect  ")  may  be  produced  on  a  salt  in 
solution  by  the  addition  to  the  solution  of  a  salt  of  the  same  base 
or  a  salt  of  the  same  acid.  The  equation  then  becomes 

a— x  =  K-x(x+p), 

p  being  the  concentration  of  the  added  ion.  K  can  only  remain 
constant  provided  x  diminishes. 

It  also  follows  that  the  degree  of  dissociation  depends  on  the  con- 
centration. If  the  latter  be  increased  n-fold,  we  have  from  the 
above  equation 

n(a-x)=Kn2x2,    or     (a -re)  -K-n-x2. 


66.]  ELECTROLYTIC  DISSOCIATION.  99 

If  n  is  >  1,  x  must  diminish,  i.e.  the  ionization  decreases  with 
increasing  concentration.  If  n  is  <1,  x  must  increase,  i.e.  the 
ionization  increases  with  the  dilution.  When  n  is  infinitely  small, 
we  have  a=x}  in  other  words,  at  infinite  dilution  the  ionization  is 
complete. 

We  are  now  able  to  give  another  definition  of  acids  and  bases 
than  that  of  §  30.  Acids  are  those  substances  which  give  H-ions 
in  aqueous  solution;  bases  under  the  same  condition  give  OH- 
ions. 

All  the  properties  of  acids,  bases  and  salts  are  closely  con- 
nected with  the  degree  of  their  ionization, — among  others  that 
which  is  indicated  by  the  rather  vague  term  strength  of  an  add 
or  base. 

As  early  as  the  eighteenth  century  it  was  observed  that  an 
acid  can  sometimes  expel  another  acid  from  its  salts.  On  adding 
hydrochloric  acid  to  sodium  carbonate,  for  instance,  sodium 
chloride  is  formed  and  carbonic  acid  given  off.  The  same  is  true 
of  bases.  When  a  solution  of  caustic  soda  is  added  to  a  solution 
of  iron  chloride,  iron  hydroxide  is  precipitated  and  sodium  chloride 
is  also  formed.  The  acid  or  base  that  can  expel  another  from  its 
salts  was  considered  by  BEKGMANN  (1735-1784)  to  be  "stronger" 
than  the  one  expelled. 

Experience  has  taught  that  those  acids  and  bases  are  strong- 
est which  are  the  most  ionized  for  the  same  dilution.  Hydro- 
chloric acid  is,  for  example,  stronger  than  hydrofluoric  acid.  At 
a  dilution  of  one  gram-molecule  per  liter  the  former  is  almost 
completely  (about  80%)  split  up  into  ions,  the  latter  only  3%. 

It  was  remarked  above  (§  30)  that  acids  turn  blue  litmus  red, 
and  bases  red  litmus  blue.  It  is  only  natural  to  seek  the  cause  of . 
these  common  properties  of  acids  on  the  one  hand  and  bases  on 
the  other  in  that  which  all  acid  solutions  have  in  common,  namely, 
hydrogen  ions,  and  in  that  which  all  solutions  of  bases  have  in 
common,  namely,  hydroxyl  ions.  The  reactions  be- 
tween acids,  bases  and  salts  in  aqueous  solu- 
tion are  almost  invariably  reactions  be- 
tween their  ions.  We  shall  explain  this  later  in  many 
instances;  the  following  example  may  suffice  for  the  present. 
When  dilute  solutions  of  a  base  and  an  acid  are  mixed,  we  have 
a  salt  solution  (§  30).  In  order  to  understand  what  reaction  has 


100  INORGANIC  CHEMISTRY.  [§  66. 

taken  place  we  must  know  that  in  dilute  solution  most  salts  are 
almost  wholly  split  up  into  ions.  Water  itself,  however,  is  split 
up  only  in  an  extremely  small  amount.  In  the  equilibrium 

H2O<=±IT+OH', 

there  is  thus  very  little  of  the  system  on  the  right-hand  side. 

The  amount  of  the  ionization  of  water  has  been  determined  in  various 
ways,  which  cannot  be  taken  up  here,  but  are  discussed  in  text-books  of 
electrochemistry.  The  results  of  the  different  methods  agree  well  and 
show  that  the  concentration  of  hydrogen,  or  hydroxyl  ions,  is  very 
nearly  1.0X10~7;  i.e.,  1  g.  H-ions  and  17  g.  OH-ions  are  contained  in 
ten  million  liters  of  water. 

Now,  when  a  base  and  an  acid  are  mixed  we  have  together  in  the 
solution  M'  +  OH'  and  A'  +  H'.  Of  these  ions  M'  and  A'  can  exist 
freely  side  by  side;  but  not  so  with  H'  +  OH',  for  these  must  unite 
to  form  water  according  to  the  above  equilibrium.  In  the  forma- 
tion of  the  salt  we  therefore  have  only  the  H'  and  OH'  ions  uniting, 
producing  undissociated  molecules  of  water. 

It  is  now  easy  to  understand  also  why  a  strong  acid  (i.e.  one 
almost  completely  ionized)  expels  a  weak  (slightly  ionized)  acid 
from  its  salts.  To  use  an  example,  suppose  we  add  to  a  liter  of 
a  sodium  fluoride  solution,  containing  one  mole  of  the  salt,  a 
similar  solution  of  hydrochloric-  acid.  In  the  mixed  solution  we 
have  the  ions 

H'  +  d'+Na' 


Since  the  equilibrium  HF<=»H'  +F  is  conditioned  on  the  presence 
of  only  3%  of  H*  Ions  and  F  ions,  there  is  a  large  excess  of  these 
ions  in  the  liquid  and  almost  all  of  them  must  unite  with  each 
other,  while  the  Na*  and  Cl'  ions  remain  free;  in  other  words, 
hydrofluoric  acid  and  sodium  chloride  (dissociated)  are  formed. 

It  also  becomes  manifest  that  the  old  notion,  once  very  gen- 
erally held,  that  the  stronger  acid  expels  the  weaker  one  from  its 
salts  completely  is  incorrect.  When  the  expelled  acid  or  base  escapes 
from  the  solution  as  a  gas  or  is  precipitated,  the  expulsion  may  in- 
deed seem  to  be  complete;  we  shall  examine  the  case  more  thor- 
oughly in  §  73. 

We  can  now  go  a  step  farther.     It  was  stated  above  that  water 


66.]  ELECTROLYTIC  DISSOCIATION.  ^';M01 

is  partially  ionized,  though  only  to  a  very  slight  extent.  Suppose 
that  we  dissolve  in  water  a  salt  of  a  strong  base  and  an  extremely 
weak  acid,  such  as  potassium  cyanide,  for  instance.  As  such  a 
salt  is  highly  ionized,  we  have  in  the  solution  the  ions 


H*  +  OH'+K' 

Since  the  acid  HCN  is  very  weak,  there  will  be  too  many  CN'  ions 
in  the  liquid  to  satisfy  the  demands  of  the  equilibrium  equation 

H'  +  CN'<=±HCN; 

hence  some  of  the  H'  ions  of  the  water  will  unite  with  CN'  ions. 
At  the  same  time,  however,  an  excess  of  OH'  ions  is  created  in 
the  liquid;  for,  inasmuch  as  potassium  hydroxide  is  a  strong  base, 
they  do  not  unite  with  the  K*  ions.  The  water,  which  originally 
reacts  neutral,  because  hydrogen  and  hydroxyl  ions  are  present  in 
equal  numbers  and  are  mutually  compensating  in  regard  to  their 
action  on  litmus,  thus  comes  to  have  an  alkaline  reaction  by  the 
solution  in  it  of  potassium  cyanide.  We  therefore  see  that  salts  of 
this  nature  are  partially  split  up  by  water  into  free  base  (K'  +  OH') 
and  free  acid  (undissociated  HCN).  This  phenomenon  is  called 
hydrolysis.  We  shall  meet  with  it  frequently  in  the  sequel. 

When  ARRHENIUS  presented  the  doctrine  of  electrolytic  dis- 
sociation in  1887,  it  met  with  much  opposition.  It  was  seen 
that  its  effect  would  be  to  produce  a  veritable  revolution  of  many 
previously  accepted  views.  Compounds  such  as  hydrochloric 
acid,  sodium  nitrate  and  others,  which  had  ever  been  considered 
as  the  most  stable,  were  to  be  supposed  according  to  the  theory 
of  ARRHENIUS  to  break  up  as  soon  as  they  dissolve  in  water.  It 
also  seemed  nonsensical  that  we  should  have  to  assume  the  exist- 
ence of  free  potassium  and  iodine  in  a  solution  of  potassium  iodide, 
for  example,  since  potassium  produces  hydrogen  and  potassium 
hydroxide  as  soon  as  it  touches  water  and  since  a  KI  solution  is 
colorless,  while  iodine  solutions  are  brown. 

So  far  as  the  first  point  is  concerned,  it  should  be  noted  that 
it  is  the  solutions  of  these  same  strongly  ionized  compounds  which 
are  chemically  the  most  active,  a  fact  which  indicates  rather  a 
loose  than  a  firm  union  of  the  constituent  atoms  in  the  molecules. 
In  regard  to  the  example  of  potassium  iodide  solution  and  other 
cases,  care  must  be  taken  not  to  confuse  atoms  and  ions.  The 


102  A 


CHEMISTRY. 


66- 


solution  of  potassium  iodide — retaining  our  illustration — contains 
neither  free  potassium  nor  free  iodine  but  ions  of  potassium  and 
ions  of  iodine.  The  atoms,  however,  must  possess  an  altogether 
different  energy  supply  than  the  ions,  whose  electric  charges  are 
very  heavy,  as  can  be  proved  by  different  methods.  It  is  this 
energy  supply  on  which  the  properties  of  bodies  depend ;  and  since 
this  is  apparently  much  different  with  the  ions  than  with  the  atoms, 
it  is  perfectly  natural  that  the  latter  should  display  other  proper- 
ties than  the  former. 

SULPHUR. 

67.  Sulphur  was  known  to  the  ancients.  It  occurs  free  in 
nature,  principally  in  the  vicinity  of  active  or  extinct  volcanoes. 
Sicily  is  its  most  important  locality,  closely  followed  by  Louisiana 
in  the  United  States,  but  large  quantities  are  also  found  in  other 
parts  of  the  United  States  and  in  Iceland,  Japan,  and  Mexico. 


FIG.  24. — DISTILLATION  OF  SULPHUR. 


It   is   separated  from  the   accompanying  rock,   or   matrix,   by 
fusion. 


67.]  SULPHUR.  103 

In  Louisiana  this  is  accomplished  by  the  FRASCH  process,  whereby 
hot  water  under  pressure  is  forced  through  pipes  sunk  through  the 
ground  to  the  sulphur  deposit,  thus  melting  the  sulphur,  which,  in  a 
molten  form,  is  forced  up  to  the  surface  by  compressed  air. 

The  crude  sulphur  thus  obtained  is  still  impure.  It  is  refined  (Fig.  24) 
by  distillation.  After  being  melted  irr  B  it  is  let  down  into  the  cast-iron 
cylinder  A,  which  is  heated  to  a  temperature  above  the  boiling-point  of 
sulphur.  The  vapor  is  conducted  into  a  large  brick  chamber,  equipped 
with  a  safety  valve  for  the  release  of  air.  If  the  distillation  is  conducted 
so  slowly  that  the  temperature  of  the  chamber  does  not  exceed  the 
boiling-point  of  sulphur,  the  latter  is  deposited  in  the  form  of  a  fine 
powder,  called  "flowers  of  sulphur" — just  as  water  vapor,  when  suddenly 
cooled  below  0°,  turns  to  snow.  Rapid  distillation,  however,  yields  a 
layer  of  liquid  sulphur  on  the  floor.  It  may  be  let  out  through  the 
opening  C  and  cast  into  slightly  conical  wooden  molds.  This  is  the  roll 
sulphur,  or  the  roll  brimstone,  of  commerce. 

Besides  occurring  in  the  free  state  sulphur  is  also  found  in 
numerous  compounds,  from  some  of  which  it  is  obtained,  e.g. 
pyrite,  or  iron  pyrites,  FeS2,  which  yields  sulphur  on  heating: 

3FeS2  =  Fe3S4+2S. 

Many  other  compounds  of  the  element  with  metals,  the  sul- 
phides, occur  in  nature,  e.g.  galenite  (lead  sulphide),  zinc 
blende  (sphalerite,  zinc  sulphide),  stibnite  (antimony  sulphide), 
cinnabar  (mercury  sulphide),  realgar  and  orpiment  (arsenic  sul- 
phides) and  chalcopyrite  (copper  pyrites,  copper  and  iron  sul- 
phide). Sulphur  also  occurs  in  the  natural  sulphates,  of  which 
gypsum  (CaS04  +  2H2O)  is  the  most  important.  It  is  also  found 
in  the  organic  world  as  a  constituent  of  the  albuminoids. 

Physical  Properties. — Sulphur  is  known  in  various  modifications. 
At  ordinary  temperatures  the  stable  form  is  a  yellow  crystalline 
solid;  melting-point,  119.25°.  A  little  above  its  melting-point 
sulphur  is  a  mobile  yellow  liquid.  With  a  continued  rise  of  tem- 
perature it  becomes  much  darker  in  color  and  very  viscid;  at 
180°  it  can  no  longer  be  poured;  at  a  higher  temperature,  espe- 
cially above  300°,  it  again  becomes  mobile,  the  dark  color  remain- 
ing; at  448°  it  boils,  producing  an  orange-colored  vapor.  At 
500°  the  vapor  is  red;  above  this  temperature  it  becomes  clearer 
again.  During  cooling  these  phenomena  reappear  in  inverse 
order.  At  -80°  sulphur  is  colorless. 


104  INORGANIC  CHEMISTRY.  [§§  67- 

Sulphur  is  insoluble  in  water  and  difficultly  soluble  in  alcohol 
and  in  ether;  it  is  easily  soluble  in  carbon  disulphide  and  in  sulphur 
monochloride,  S2C12.  100  parts  CS2  dissolve  46  parts  S  at  22°. 

The  molecular  weight  of  this  element,  more  than  that  of  any 
other,  depends  on  the  temperature.  Below  the  boiling-point  the 
molecular  formula  is  Sg,  according  to  the  determination  of  the 
boiling-point  elevation  in  carbon  disulphide  (boiling-point  46°) 
and  the  freezing-point  depression  of  fused  naphthalene  (melting- 
point  80°).  In  the  gaseous  state  the  density  (air=  1)  varies  from 
7.937  at  467.9°  to  2.23  at  860°  and  then  remains  constant  even 
as  high  as  1800°,  indicating  that  at  the  lowest  temperatures 
sulphur  vapor  consists  of  Sg  molecules,  and  above  860°  of  only 
82  molecules. 

Above  1800°  the  molecule  82  begins  to  dissociate  into  its 
atoms;  at  2000°  and  0.5  atmosphere  pressure  the  dissociation  has 
reached  about  45%,  according  to  an  investigation  of  NERNST. 

68.  Allotropic  Modifications. — At  least  four  solid  forms  are 
known,  while  in  the  liquid  state  there  are  two  more.  The  solid 
allotropic  conditions  can  be  divided  into  crystallized  and  amor- 
phous. As  for  the  former,  sulphur  is  dimorphic,  forming  rhombic 
as  well  as  monoclinic  crystals.  The  former  are  transformed  into 
the  latter  on  heating  (§  70). 

Rhombic  sulphur  can  be  obtained  in  beautiful  crystals  by  allowing  a 
solution  of  sulphur  in  carbon  disulphide  or  chloroform  to  slowly  evaporate. 
Monoclinic  sulphur  is  easily  obtained  in  the  following  manner:  Some 
sulphur  is  fused  in  a  large  crucible  and  allowed  to  cool  slowly  until  a  crust 
forms  on  the  surface.  The  crust  is  then  broken  through  and  the  liquid 
sulphur  poured  out;  the  sides  of  the  crucible  are  found  to  be  covered 
with  long,  yellow  transparent  needles.  In  the  course  of  a  few  hours  these 
become  opaque  and  brittle,  however,  and  crumble  at  the  slightest  touch 
to  a  powder,  which  is  found  to  consist  of  rhombic  crystals  (cf.  §  71). 

Amorphous  sulphur  may  be  either  soft  and  soluble  in  carbon 
disulphide  or  powdery  and  insouble  in  this  liquid. 

The  soluble  kind  results  from  the  decomposition  of  certain  sulphur 
compounds.  When  hydrogen  sulphide  water  is  exposed  to  the  air,  sul- 
phur slowly  separates  in  the  form  of  a  white  powder.  The  polysulphides 
(CaSn,  K2Sn,  etc.)  yield,  when  decomposed  by  an  acid,  a  cloudy  milk-like 
liquid,  which  is  found  to  contain  extremely  fine  particles  of  amorphous 


68.]  SULPHUR.  105 

sulphur.     In  either  case  there  is  always  formed  in  addition  to  the  soluble 
variety  some  insoluble  (in  CS2)  sulphur. 

The  insoluble  form  may  be  best  prepared  by  heating  sulphur  to 
near  its  boiling  point  and  then  pouring  it  in  a  fine  stream  into  cold 
water;  thereby  a  semi-solid  plastic  modification  is  formed,  which  becomes 
brittle  after  a  time.  By  extraction  with  carbon  disulphide  the  soluble 
modification  is  removed  and  a  yellow  powder  remains,  which  is  the 
amorphous  modification,  insoluble  in  that  liquid.  The  relative  quantity . 
of  this  latter  modification  depends  only  on  the  temperature  at  which 
the  sulphur  was  heated.  The  higher  the  temperature,  the  greater  the 
yield;  on  heating  at  440°,  for  example,  the  yield  is  30.3%.  It  is  a  very 
curious  fact  that  amorphous  sulphur  is  not  formed  if  a  few  bubbles  of 
ammonia  gas  or  of  carbon  dioxide  are  first  passed  into  the  heated  sul- 
phur and  that,  on  the  contrary,  the  introduction  of  air  will  restore  the 
ability  to  form  amorphous  sulphur.  Probably  traces  of  sulphur  dioxide 
are  necessary  for  the  formation  of  the  amorphous  state.  Further,  it  has 
been  proved  that  molten  sulphur,  no  matter  whether  it  can  produce  the 
amorphous  modification  by  rapid  cooling  or  not,  has  in  both  cases  the 
same  physical  properties,  such  as  specific  gravity,  boiling-point,  solu- 
bility, etc. 

In  order  to  explain  this  fact,  the  conduct  of  sulphur  on  heating  must 
be  considered.  As  has  already  been  stated,  sulphur  when  heated  above 
its  melting-point  is  at  first  a  mobile  liquid.  When  the  temperature 
reaches  160°  the  liquid  very  soon  becomes  viscid.  If  the  temperature 
is  maintained  for  some  time  at  158-160°,  the  molten  mass  separates  into 
two  liquid  layers,  a  mobile  one  and  a  viscid  one.  In  other  words,  sul- 
phur can  form  two  liquid  modifications,  which  are  only  partially 
miscible.  At  every  temperature  an  equilibrium  establishes  itself  in  the 
molten  sulphur  between  these  modifications.  When  rapidly  cooled  the 
viscid  form  continues  to  exist,  for  the  reason  that  its  transition  velocity 
is  strongly  diminished,  and  gives  the  amorphous  modification  on  solidi- 
fying. When  cooled  slowly  the  viscid  modification  changes  gradually 
into  the  mobile  liquid  form  and  the  resulting  solid  does  not  contain 
amorphous  sulphur.  That  amorphous  sulphur  is  not  obtained  by  rapidly 
cooling  molten  sulphur  which  is  completely  free  from  sulphur  dioxide, 
must  probably  be  attributed  to  a  very  strong  catalytic  retardation  of 
the  transition  of  the  viscid  form  into  the  mobile  one  by  traces  of  sulphur 
dioxide. 

In  the  light  of  the  above  facts  concerning  the  behavior  of  sulphur  we 
can  understand  why  its  melting-point  is  dependent  on  its  previous 
history.  Sulphur  that  has  been  heated  melts  lower  than  otherwise. 
This  is  due  to  the  presence  of  amorphous  sulphur  and  the  consequent 
lowering  of  the  melting-point,  the  same  as  by  a  foreign  substance. 

The  molecular  weight  of  the  amorphous  insoluble  sulphur  is  also  S8. 


106  INORGANIC  CHEMISTRY.  [§§e9- 

69.  Chemical    Properties.  —  Sulphur    combines    directly    with 
many  elements,  not  only  metals  but  also  metalloids.     It  has  been 
already  stated  (§  10)  that  it  burns  with  a  blue  flame  when  heated 
in  air  or  in  oxygen.     The  halogens  and  hydrogen  unite  with  it 
directly.     Powdered  iron  and  sulphur,  when  mixed  and  heated, 
combine  energetically,  producing  great  heat  (§  20).     Copper  takes 
fire  in  the  vapor  of  boiling  sulphur.     When  mercury  and  sulphur 
are  rubbed  together  in  a  mortar,  black  mercuric  sulphide,  HgS,  is 
formed.     The  sulphur  compounds  of  the  metals  are  called  sul- 
phides. 

THE  TRANSITION  POINT. 

70.  As  stated  in   §  68  sulphur  can  crystallize  in  two  modi- 
fications, rhombic  .and  monoclinic.    These   modifications   can   be 
readily  transformed  into  one  another.      The  peculiar  phenomena 
connected  with  this  transition  deserve  a  closer  study.     At  ordi- 
nary temperatures  sulphur  is  rhombic   and  remains  so   till  the 
temperature  95.4°  is  reached,  above  which  there  begins  a  slow  but 
complete  transformation  into  the  monoclinic  variety.     Inversely, 
when  the  monoclinic  modification  is  subjected  to  a  temperature 
below  95.4°,  a  complete  change  into  the  rhombic  form  occurs. 
At   the   temperature   named   the   two   modifications   are   equally 
stable  and  can  exist  side  by  side  in  any  proportions  for  an  indefinite 
period;   above  it  only  the  monoclinic,  below  it  only  the  rhombic, 
form  can  exist  permanently.     Such  phenomena  are  not  infrequent. 
The  temperature  at  which  the  one  system  passes  into  the  other 
is  called  the  transition  point,  also  point  of  inversion.    This  transi- 
tion point  possesses  great  analogy  with  the  melting-point.     Just 
as  ice,  for  example,  is  changed  into  water  above  0°  and  water  into 
ice  below  0°,  so  in  a  system  of  substances  possessing  a  transition 
point  only  one  system  is  stable  below  that  point,  above  it  only 
the  other. 

^The  theoretical  explanation  of  both  phenomena  is  exactly  the 
same.  Let  us  consider  a  body,  ice  for  example,  at  temperatures 
slightly  below  its  melting-point,  and  represent  graphically  in  the 
diagram  OTP  (Fig.  25)  the  values  of  the  vapor  tension  corre- 
sponding to  different  temperatures.  The  result  is  the  line  marked 
ice  in  the  figure.  This  vapor-tension  curve,  if  prolonged  through 
and  beyond  the  melting-point,  is  found  to  bend  sharply  at  the 


70.] 


THE  TRANSITION  POINT. 


107 


latter  and  take  a  new  direction.  This  deflection  is  very  slight  in 
the  case  of  ice  and  water;  it  can  be  nevertheless  experimentally 
detected;  it  is  much  more  evident  with  benzene  and  many  other 
substances.  By  carefully  cooling  water  it  can  be  made  to  remain 
liquid  even  under  0°;  such  a  liquid  is  said  to  be  supercooled. 
The  vapor  tension  of  this  supercooled  water  is  greater  than  that  of 
ice  at  the  same  temperature  and  the  curve  representing  the  former 
is  but  a  continuation  of  the  vapor-tension  curve  for  water.  Since 
the  vapor  tension  of  supercooled  water  is  greater  than  that  of  ice, 
water  at  temperatures  below  0°  must,  according  to  previous  con- 
clusions (§  43,  3),  pass  into  ice  when  the  two  are  in  contact.  How- 
ever, the  vapor  tension  of  water  at  a  temperature  slightly  above 


o  o°  T 

FIG.  25. 


FIG.  26. 


0°  will  be  less  than  that  of  ice  and  we  shall  have  the  ice  trans- 
formed into  water.  It  is  therefore  evident  that  both  above 
and  below  the  melting-point  one  of  the  systems  will  necessarily 
disappear. 

Exactly  the  same  explanation  can  be  offered  for  the  transition 
point.  Below  95.4°  the  vapor  tension  of  rhombic  sulphur  is  less 
than  that  of  monoclinic  sulphur;  above,  the  vapor  tension  of  the 
rhombic  variety  exceeds  that  of  the  monoclinic.  There  is  therefore 
a  complete  transformation  from  one  system  to  the  other  when 
the  temperature  is  other  than  95.4°,  for  the  same  reason  as  in  the 
case  of  the  melting-point;  moreover,  just  as  ice  and  water  under 
ordinary  pressure  can  exist  side  by  side  indefinitely  only  at  0°,  so 
both  modifications  of  sulphur  are  coexistent  only  at  95. 4°,  since 
only  then  is  the  vapor  tension  the  same  for  both  systems 
(Fig.  26). 

Of  the  various  methods  for  the  determination  of  the  transition 


108  INORGANIC    CHEMISTRY.  [§§68^ 

point  a  convenient  one  is  the  dilatometric  method.  It  is  based  on 
the  change  of  volume  (specific  gravity)  which  a  body  usually 
undergoes  on  passing  through  the  transition  point.  In  measuring 
this  a  dilatometer  is  used,  an  instrument  which  may  be 
compared  to  a  thermometer  of  very  large  dimensions.  After 
rhombic  sulphur,  for  example,  has  been  placed  in  the  dilatometer 
the  latter  is  filled  with  a  chemically  indifferent  liquid  (kerosene, 
linseed  oil)  and  put  in  a  large  water  bath;  the  temperature  is 
then  slowly  raised.  Below  the  transition  point  the  volume  is 
seen  to  slowly  and  steadily  increase  with  the  temperature  on 
account  of  expansion;  as  soon  as  the  temperature  gets  a  trifle 
above  95.4°,  however,  a  marked  increase  of  volume  is  observed, 
even  if  the  temperature  be  maintained  constant;  thereupon 
expansion  again  proceeds  gradually,  as  before,  if  the  tem- 
perature is  allowed  to  rise.  The  marked  change  of  volume 
indicates  the  transition  of  the  rhombic  sulphur  into  the  mono- 
clinic  modification. 


"STABLE,"   "METASTABLE,"   AND   "LABILE." 

These  terms  are  coming  to  be  so  frequently  used  in  chemistry 
that  they  need  to  be  distinctively  denned.  They  are  borrowed 
from  mechanics,  for  which  reason  it  is  desirable  that  they  be 
employed  in  chemistry  in  the  same  sense  as  in  mechanics.  In  the 
latter  an  equilibrium  is  called  labile  (apt  to  slip)  when  the  slightest 
displacement  suffices  to  transpose  the  body  into  a  new  position  of 
equilibrium.  An  example  is  afforded  by  a  cone  standing  on  its 
apex.  It  cannot  recover  from  even  the  slightest  disturbance, 
but  gets  further  and  further  from  the  vertical  position  and  finally 
tumbles  over.  A  labile  condition  is  thus  really  a  limiting  case 
which  cannot  actually  be  realized;  not  even  for  the  cone,  though 
its  apex  were  a  mathematical  point  resting  on  an  absolutely  hard 
surface. 

All  actually  occurring  equilibria  are  stable;  but  there  can  be 
different  degrees  of  stability.  When  a  material  cone  is  stood  on 
its  apex  its  equilibrium  has  very  little  stability.  On  the  contrary 


71.]  THE  PHASE  RULE  OF  GIBBS.  109 

a  beam  resting  on  the  ground  with  its  largest  surface  down 
represents  a  very  stable  equilibrium.  However,  if  the  beam 
is  stood  up  on  end,  its  equilibrium  becomes  less  stable.  Like 
the  cone  resting  on  its  apex,  the  beam  will  have  a  tendency 
to  go  over  into  a  more  stable  condition.  In  mechanics  there 
is  no  need  of  giving  such  conditions  a  special  name,  but  in 
thermodynamics  and  chemistry  they  call  for  special  designa- 
tion, and  the  term  applied  to  them  is  metastable.  We  have  an 
example  in  undercooled  water,  something  that  can  exist,  but 
has  the  tendency  to  go  over  into  a  more  stable  condition, 
namely,  into  ice.  Therefore  undercooled  water  is  said  to  be 
metastable. 

It  follows  from  the  above  that  expressions  such  as,  "  labile 
compounds "  (e.g.,  for  C102),  or  "  the  substance  exists  in  a 
labile  condition,"  are  to  be  avoided.  The  word  "  labile " 
should  be  replaced  by  "  metastable."  Strictly  labile  condi- 
tions are  impossible;  nevertheless  they  may  possess  great 
theoretical  interest,  such,  for  instance,  as  the  case  of  the 
continuous  transformation  of  liquid  into  gas  below  the 
critical  temperature,  which,  though  it  cannot  be  carried 
out,  has  led  to  very  valuable  theoretical  considerations  in  the 
hands  of  VAN  DER  WAALS  and  others,  as  may  be  seen  in  the 
larger  text-books  of  physics. 


THE  PHASE  RULE   OF   GIBBS. 

71.  The  phase  rule  treats  of  the  equilibrium  in  heterogeneous 
systems,  i.e.  systems  that  can  be  separated  mechanically  into 
unlike  parts.  A  saturated  salt  solution  in  contact  with  solid 
salt  is  a  heterogeneous  system,  for  it  consists  of  solid  salt,  the 
solution  and  vapor;  that  is,  of  three  parts,  mechanically  sepa- 
rable. Each  of  these  parts  in .  itself  is  homogeneous, 
i.e.  each  part  has  the  same  composition  throughout.  A  gas 
mixture  is  always  homogeneous,  as  is  also  a  solution.  These 
homogeneous  parts,  separated  by  limiting  surfaces  and  of 
which  a  heterogeneous  system  is  made  up,  are  called  by  GIBBS 


110  INORGANIC  CHEMISTRY.  [§§71- 

phases.     Water  and  its  vapor  constitute  two  phases;   ice,  water 
and  steam  three  phases. 

A  heterogeneous  system  can  never  have  more  than  one 
gaseous  phase,  because  all  gases  are  miscible  in  all  proportions* 
it  may,  however,  consist  of  different  liquid  phases,  in  case  it 
contains  immiscible  liquids.  The  number  of  these  liquid  phases 
is  seldom  more  than  two;  that  of  the  solid  phases  is  un- 
limited. 

A  further  conception,  introduced  by  GIBBS,  is  that  of  the 
components  of  a  system.  If  the  system  is  composed  of  only 
one  element,  then  this  element  is  the  only  component.  Systems 
made  up  of  one  compound  have  in  most  cases  this  compound 
as  the  only  component.  A  system  consisting  of  molten  and 
gaseous  sulphur,  or  of  water  and  steam,  has  but  one  component. 
In  this  case  all  phases  have  the  same  composition.  There  are 
systems,  however,  in  which  this  is  not  the  case;  viz.  systems 
that  are  made  up  of  more  than  one  component.  We  select  as 
the  components  those  compounds  of  which  the  smallest  number 
is  necessary  to  form  the  different  phases.  The  choice  of  such 
compounds  may  be  somewhat  arbitrary  but  their  number  is 
always  fully  denned. 

Let  us  consider,  for  example,  the  system  Glauber's  salt- 
water. This  salt  has  the  composition,  Na2S04-  lOH^O.  In 
order  to  determine  the  composition  of  the  phases  that  are  possible 
here  (solid  salt,  solution,  vapor)  it  is  best  to  choose  Na2SO4  and 
H2O  as  components.  We  might  indeed  take  Glauber's  salt  itself 
as  one  of  the  components;  but  then,  in  case  the  solid  phase  was 
the  anhydrous  salt,  it  would  be  necessary  to  regard  water  as  a 
negative  part  of  it,  which  is  undesirable.  Sulphuric  acid  and 
sodium  hydroxide  are  not  components,  because  they  do  not  occur 
independently  in  any  phase,  neither  are  they  found  in  any  other 
relation  in  the  phases  than  as  a  part  of  the  salt  itself.  It  is  a 
property  of  the  components  that  they  can  occur  in  some  of  the 
phases  in  varying  proportions  (e.g.  in  saturated  and  unsaturated 
solutions). 

Let  us  now  take,  for  example,  a  saturated  solution  of  salt  and 
water  in  a  vessel  that  is  closed  with  a  movable  piston.  Under 
this  solution  let  there  be  a  little  solid  salt,  above  it  the  vapor  of 


71.]  THE   PHASE  RULE   OF   GIBBS.  HI 

the  solution.  The  system  consists  manifestly  of  two  substances 
and  three  phases. 

So  long  as  the  temperature  remains  constant,  the  vapor  of  the 
salt  solution  possesses  a  definite  tension.  If  we  increase  the 
volume  by  raising  the  piston,  a  definite  amount  of  water  will  evap- 
orate; since  the  solution  was  saturated,  the  result  will  be  that  a 
little  salt  will  be  deposited;  in  the  end  the  quantities  of  vapor, 
solution  and  salt  will  therefore  have  altered,  but  the  composition 
of  each  phase  will  remain  the  same.  The  tension,  and  hence  also 
the  concentration,  of  the  vapor  remain  unchanged,  since  the 
temperature  is  constant;  there  is  likewise  no  change  in  the  con- 
centration of  the  salt  solution.  The  same  is  true  in  case  the 
volume  be  diminished.  It  therefore  follows  that  the  equilibrium 
of  such  a  system  is  independent  of  the  quantities  of  the  various 
phases.  It  is  dependent  only  on  the  temperature  chosen;  if 
this  is  constant,  the  whole  system  is  defined.  Or,  if  we  should 
select  an  arbitrary  value  for  the  composition,  the  temperature 
and  pressure  would  be  fully  defined.  It  is  therefore  evident 
that  the  system  is  completely  defined  as  soon  as  one  of  these 
magnitudes  is  arbitrarily  chosen.  The  system  has  only  one 
degree  of  freedom.  Such  an  equilibrium  possesses  the  following 
characteristic:  At  a  given  constant  temperature  the  vapor  pres- 
sure is  definite.  Under  an  even  slightly  greater  or  smaller 
pressure  one  of  the  phases  will  gradually  and  completely  dis- 
appear, provided  the  temperature  remains  constant.  On  in- 
creasing the  pressure  the  gaseous  phase  wholly  condenses,  so  that 
only  solution  and  salt  remain.  A  decrease  of  pressure  results  in 
the  complete  evaporation  of  the  solution,  vapor  and  salt  only 
being  left. 

The  same  is  true  when  the  pressure  remains  constant  and  the 
temperature  varies. 

An  entirely  different  behavior  is  shown  by  a  system  made  up 
of  an  unsaturated  salt  solution  and  its  vapor.  At  a  constant 
temperature  and  a  definite  position  of  the  piston  the  vapor  tension 
has  a  definite  value,  as  in  the  former  case.  If,  however,  the 
volume  of  vapor  be  changed,  the  tension  will  correspondingly 
vary,  for,  if  the  volume  be  increased,  for  example,  more  water 
will  evaporate,  the  solution  will  become  more  concentrated  and 
the  vapor  tension  of  course  lessen.  Therefore  for  every  definite 


112 


INORGANIC    CHEMISTRY. 


[§  71- 


temperature  there  are  not  simply  one  but  infinitely  many  pres- 
sures under  which  this  system  can  be  in  equilibrium.  The  result 
is  that  the  slightest  change  of  volume  or  pressure  does  not 
necessitate  the  disappearance  of  one  of  the  phases.  Two 
magnitudes  may  be  chosen  arbitrarily  before  the  system  is  fully 
denned;  it  has  two  degress  of  freedom.  It  is  evident  in  this 
example  that  the  number  of  degrees  of  freedom  increases  by  one 
when  the  number  of  phases  decreases  by  one. 

The  phase  rule  expresses  a  relation  between  the  numbers  of 
the  components  S,  the  phases  P,  and  the  degrees  of  freedom  F. 
It  is  of  the  following  form: 


or,  in  words,  The  sum  of  the  number  of  the  degrees  of  freedom  and 
the  number  of  the  phases  of  a  system  exceeds  the  number  of  com- 
ponents by  two. 

Let  us  apply  the  phase  rule  in  the  first  place  to  water,  a  system 
of  one  component;  the  sum  of  the  degrees  of  freedom  and  the 
phases  must  therefore  be  three. 

In  the  following  graphic  Representation,  Fig.  27,  the  tempera- 


Solid 


Liquid 


! 
Gaseous 


0° 
FiG.  27. 


95.°4  120° 

FIG.  28. 


tures,  t,  are  plotted  as  abscissie,  the  pressures,  P,  as  ordinates. 
Let  us  first  consider  liquid  water  above  0°.  The  number  of  the 
phases  is  two  (liquid  and  vapor) ;  the  system  has  therefore  only 
one  degree  of  freedom,  or,  as  we  say,  it  is  univariant.  To  every 
temperature  there  corresponds  a  definite  vapor  tension.  The 
ordinates  of  every  point  in  the  line  OB  indicate  these  vapor  ten- 


71.]  THE    PHASE   RULE    OF  GIBBS.  113 

sions.  If  the  pressure  at  a  certain  temperature  were  greater  than 
that  indicated  by  the  ordinate,  the  gaseous  phase  would  com- 
pletely disappear.  The  line  OB  therefore  represents  the  boundary 
between  the  liquid  and  gaseous  phases  for  the  various  temperatures 
and  pressures.  Every  point  in  the  area  COB  represents  the  liquid, 
every  point  in  A  OB  the  gaseous,  phase.  Only  the  points  of  the 
line  OB  indicate  the  temperatures  and  corresponding  pressures, 
at  which  both  phases  are  coexistent.  The  line  OB  therefore 
ends  on  one  side  at  0°;  its  other  end  is  at  the  critical  temperature, 
since  at  this  point  vapor  and  liquid  become  identical. — Let  us  now 
allow  the  temperature  to  fall  below  0°.  The  liquid  phase  dis- 
appears and  ice  takes  its  place.  The  system  remains  unvariant, 
however,  for  the  number  of  phases  is  unchanged.  The  ordinates 
of  the  points  on  the  line  OA  again  give  the  vapor  tensions  of  ice 
for  different  temperatures.  For  the  same  reason  as  above  OA 
is  the  boundary  line  between  the  solid  and  gaseous  phases.  Only 
along  this  line  are  the  two  coexistent.  The  line  OA  extends  to 
the  absolute  zero,  since  the  gaseous  phase  then  disappears. 

The  melting-point  of  ice  depends  somewhat  on  the  pressure, 
being  lowered  by  increasing  pressure  0.0075°  per  atmosphere. 
Both  phases,  ice  and  water,  will  therefore  be  coexistent  along  the 
line  OC,  which  shows  a  very  considerable  rise  of  pressure  for  a 
very  slight  fall  of  temperature.  In  this  case  also  a  change  of 
pressure  at  a  constant  temperature,  or  the  reverse,  involves  the 
complete  disappearance  of  one  of  the  phases.  The  line  OC  will 
end  at  a  point  where  the  liquid  and  solid  phases  become  identical, 
i.e.  where  the  whole  system  turns  to  a  homogeneous  amorphous 
mass.  The  location  of  this  point  has  not  yet  been  ascertained. 

The  point  0  (about  +0°.01),  the  melting-point  of  ice  at  the 
pressure  of  the  vapor,  is,  according  to  the  above  mode  of  repre- 
sentation, the  point  of  intersection  of  the  three  lines  which  sepa- 
rate the  phases  and  along  which  two  phases  are  coexistent.  It  is 
called  a  triple  point.  Only  in  this  point  is  it  possible  for  the 
three  phases  to  exist  side  by  side;  it  is  the  common  point  of  the 
areas  which  represent  regions  of  the  three  phases.  When  three 
phases  are  coexistent  the  system  has  no  degree  of  freedom;  it 
has  become  non-variant. 

In  the  case  of  sulphur  we  have  one  substance  and  four  possible 
phases:  rhombic,  monoclinic,  liquid,  gaseous.  Fig.  27  makes 


114  INORGANIC  CHEMISTRY.  [§§71- 

plain  the  relation  between  these  phases.  Below  95.4°  sulphur  is 
rhombic;  the  two  phases  are  rhombic  sulphur  and  vapor.  The 
line  OA  forms  the  boundary  between  the  two  regions.  At  95.4° 
the  rhombic  phase  passes  into  the  monoclinic  phase.  The  ordi- 
nates  of  the  line  OB  represent  the  vapor  pressure  of  monoclinic 
sulphur  at  the  temperatures  95.4°-120°.  The  two  crystallizablc 
phases  can  exist  side  by  side  at  the  point  0  (the  transition  point). 
According  to  researches  by  REICHER  this  transition  point  depends 
on  the  pressure;  an  increase  of  pressure  of  one  atmosphere  raises 
it  about  0.05°.  The  boundary  between  the  crystallized  phases  is 
therefore  furnished  by  a  line  OC,  which  shows  that  a  very  slight 
rise  of  temperature  is  followed  by  a  very  considerable  increase  of 
pressure.  At  0  we  have  therefore  a  triple  point,  i.e.  a  point  com- 
mon to  both  crystallized  phases  and  the  gaseous  phase.  At  B,  the 
melting-point  of  monoclinic  sulphur,  there  is  a  second  triple  point, 
which  is  wholly  analogous  to  the  melting-point  of  ice.  Finally,  it 
should  also  be  noted  that  the  line  BC',  which  separates  the  liquid 
and  the  solid  phases,  must  indicate  a  rise  of  melting-point  for  an 
increase  of  pressure,  since  sulphur  melts  higher  the  greater  the 
pressure.  The  lines  OC  and  EC'  are  not  parallel  but  intersect, 
according  to  TAMMANN'S  experiments,  at  151°  and  1281  atmos- 
pheres. As  the  sum  of  the  phases  and  degrees  of  freedom  is  also 
three  with  sulphur,  the  phase  rule  indicates  that  all  four  phases 
cannot  exist  in  the  presence  of  each  other  at  the  same  time,  not 
even  when  the  system  has  become  non-variant. 

At  the  triple  point  neither  the  temperature  nor  the  pressure 
can  be  changed  without  altering  the  kind  of  equilibrium.  Here 
the  system  is  non-variant.  Along  the  lines  OA,  OB  and  OC  it 
is  univariant.  When  the  state  of  the  system  is  represented  by  a 
point  within  one  of  the  areas  it  is  divariant,  consisting  then  of 
only  one  phase.  In  the  succeeding  chapters  we  shall  have 
occasion  to  concern  ourselves  with  systems  of  more  than  one 
component. 


72.]  HYDROGEN  SULPHIDE.  115 

HYDROGEN  SULPHIDE,  SULPHURETTED  HYDROGEN,  H2S. 

72.  This  gas  occurs  in  nature  chiefly  in  volcanic  regions.  Cer- 
tain mineral  waters,  especially  the  so-called  "sulphur  springs/' 
contain  it.  It  is  also  found  as  a  putrefactive  product  of  organic 
bodies. 

Hydrogen  sulphide  can  be  obtained  from  its  elements  by 
synthesis.  They  unite  almost  completely  when  heated  together 
for  a  long  time  (about  168  hours)  at  310°. 

It  can  also  be  obtained  by  the  action  of  hydrogen  on  sulphur 
compounds,  as  well  as  by  the  action  of  sulphur  on  compounds  of 
hydrogen;  the  reduction  of  silver  sulphide,  Ag2S,  with  hydrogen 
at  high  temperatures  illustrates  the  former  case,  while  the  boiling 
of  turpentine  oil  with  sulphur  is  an  example  of  the  latter. 

None  of  the  above  methods  is  adapted  to  the  preparation  of 
the  gas  in  the  laboratory.  For  this  purpose  the  interaction  of  a 
sulphide  with  a  hydrogen  compound  is  employed,  iron  sulphide 
and  dilute  acids  being  generally  used: 

FeS + 2HC1  =  FeClp + H2S. 

In  order  to  have  sulphuretted  hydrogen  always  at  hand,  it  being  in  con- 
stant demand  in  analytical  work  (cf.  §  73),  a  very  convenient  apparatus 
was  devised  by  KIPP,  which  can  be  used  for  the  generation  (at  ordinary 
temperatures)  of  other  gases  as  well.  Its  construction  is  shown  in  the 
figure  (see  next  page) . 

The  lower  globe  is  joined  to  the  basal  portion  by  a  narrow  neck,  while 
the  upper  globe  tapers  into  a  long  tube,  which  fits  tightly  into  the  lower 
globe  and  extends  nearly  to  the  bottom  of  the  generator  without  com- 
pletely filling  the  neck.  The  iron  sulphide  is  put  into  the  middle  portion 
and  the  dilute  acid  is  poured  into  the  upper  portion,  the  stopcock  remain- 
ing open.  As  soon  as  the  basal  part  is  filled  with  the  acid  the  cock  is 
closed  and  the  top  part  is  half  filled  with  more  acid.  When  the  cock  is 
opened  the  liquid  sinks  in  the  top  part  and  rises  into  the  middle  portion, 
where  it  reacts  with  the  iron  sulphide  to  produce  hydrogen  sulphide, 
which  escapes  through  the  cock.  On  closing  the  latter  the  gas  continues 
to  be  evolved  till  it  forces  the  liquid  back  out  of  the  part  containing  the 
iron  sulphide.  The  reaction  thus  ceases  automatically  and  the  generator 
is  ready  at  any  time  to  supply  new  quantities  of  gas  on  opening  the  cock, 
till  either  acid  or  sulphide  is  exhausted.  The  spent  acid  can  be  let  out 
through  a  stoppered  opening  near  the  bottom. 

On  account  of  the  free  iron  usually  present  in  iron  sulphide,  the  gas 


116  INORGANIC  CHEMISTRY.  [§§72- 

prepared  in  this  manner  contains  some  hydrogen.  Perfectly  pure 
hydrogen  sulphide  is  obtained  by  warming  antimony  sulphide,  Sb2S3, 
with  concentrated  hydrochloric  acid- 


FIG.  29.  —  KIPP  GENERATOR. 

Physical  Properties.  —  Hydrogen  sulphide  is  a  colorless  gas  of 
disagreeable  odor,  when  diluted  reminding  one  of  rotten  eggs. 
Under  a  pressure  of  about  17  atmospheres  it  becomes  liquid  at 
ordinary  temperatures;  liquid  hydrogen  sulphide  boils  at  —61.8° 
and  freezes  at  —85°.  1  1.  H2S  gas  weighs  1.5392  g.  at  0°  and 
760  mm.  pressure.  The  gas  is  rather  soluble  in  water,  1  vol. 
water  dissolving  4.37  vols.  K2S  at  0°  ("  hydrogen  sulphide 
water  "). 

Chemical  Properties.  --Hydrogen  sulphide  is  combustible  and 
yields  on  combustion  either  sulphur  dioxide  and  water  or  water 
and  sulphur,  according  to  the  air  supply: 


In  aqueous  solution  it  is  slowly  oxidized  by  the  oxygen  of  the 
air,  sulphur  being  set  free;  this  decomposition  is  aided  by  light. 
In  order  to  preserve  hydrogen  sulphide  water,  it  must  be  pre- 
pared from  boiled  (air-free)  water  and  put  into  a  dark  bottle, 
filled  entirely  and  closed  air-tight.  The  latter  condition  is  best 
met  by  placing  the  bottle,  stopper  downwards,  in  a  glass  of  water. 
It  is  poisonous  ;  as  an  antidote  very  dilute  chlorine  may  be  inhaled. 


73.]  HYDROGEN  SULPHIDE.  117 

Hydrogen  sulphide  is  a  powerful  reducing-agent.  Bromine 
water  and  iodine  solution  are  decolorized  by  it  with  separation  of 
sulphur  (§§  45  and  48). 

Various  oxygen  compounds  are  transformed  by  hydrogen  sul- 
phide into  compounds  with  less  oxygen,  e.g.  chromic  acid  is 
reduced  in  acid  solution  to  a  chromic  salt  (§  295).  Fuming  nitric 
acid  acts  so  vigorously  that  a  slight  explosion  occurs.  When 
hydrogen  sulphide  is  passed  over  lead  dioxide,  the  gas  ignites, 
while  the  oxide  is  reduced.  Concentrated  sulphuric  acid,  H2SO4, 
is  also  reduced;  hence  it  cannot  be  used  for  drying  the  gas. 

Hydrogen  sulphide  possesses  the  character  of  a  weak  acid; 
when  it  is  passed  over  zinc,  copper,  tin  or  lead,  the  corresponding 
sulphides  are  formed  and  hydrogen  is  set  free. 

Composition  of  Hydrogen  Sulphide. — When  a  bit  of  tin  is  heated 
in  dry  hydrogen  sulphide — in  a  tube  over  mercury — tin  sulphide 
and  hydrogen  are  formed.  After  cooling  it  is  seen  that  the  volume 
of  hydrogen  is  just  as  great  as  that  of  the  hydrogen  sulphide. 
The  same  result  is  obtained  when  a  platinum  wire  is  heated  to 
redness  (by  an  electric  current)  in  the  dry  gas,  causing  the  latter  to 
break  up  into  its  elements.  Since 
the  hydrogen  molecule  is  H2,  there 
must  be  two  atoms  of  hydrogen  in 
the  hydrogen  sulphide  molecule. 
Now  the  specific  gravity  of  hydro- 
gen sulphide  has  been  found  to  be 

1.1912    for  air=l,  or   17.15   for 
/•v     «  A   mi  i       i     ii  FIG.  30. — DECOMPOSITION  OF  H,S. 

O  =  16.  The  gram-molecule  there- 
fore weighs  34.30  g.,  and,  as  it  contains  2  g.  hydrogen,  there 
remains  for  sulphur  32.3  g.  This  figure  is  very  close  to  the  atomic 
weight  of  sulphur,  hence  there  can  only  be  one  atom  of  sulphur 
present  in  the  molecule  of  hydrogen  sulphide.  We  thus  conclude 
that  the  formula  is  H2S. 

73.  Use  of  Hydrogen  Sulphide  in  Analysis. — Hydrogen  sulphide 
finds  extensive  use  in  qualitative  analysis.  A  large  number  of 
metals  are  precipitated  by  it  from  acid  solutions  as  sulphides,  viz., 
gold,  platinum,  arsenic,  antimony,  tin,  silver,  mercury,  lead,  bis- 
muth, copper  and  cadmium,  and  also  certain  rare  elements.  Some 
of  these  sulphides  have  a  characteristic  color,  e.g.  the  orange- 
red  antimony  sulphide,  Sb2S3,  the  yellow  cadmium  sulphide, 


118  INORGANIC  CHEMISTRY.  [§  73 

CdS,  the  brown  stannous  sulphide,  SnS,  the  yellow  stannic  sulphide, 
SnS2,  and  the  yellow  sulphides  of  arsenic,  As2S3  and  As2Ss.  The 
rest  of  the  sulphides  named  are  black.  Other  metals,  such  as 
nickel,  cobalt,  iron,  manganese,  zinc,  chromium,  aluminium,  etc., 
are  not  precipitated  by  hydrogen  sulphide  from  acid  solution .  but 
are  precipitated  by  ammonium  sulphide.  Still  other  metals,  such 
as  barium,  strontium,  calcium,  magnesium,  and  the  alkalies,  are 
not  precipitated  from  their  solutions  even  by  ammonium  sul- 
phide, so  that  we  therefore  possess  in  sulphuretted  hydrogen  and 
its  ammonium  compound  a  means  of  separating  these  elements. 
An  answer  to  the  question,  why  some  elements  are  precipitated 
from  acid  solution  by  hydrogen  sulphide  and  others  are  not,  is 
furnished  by  the  ionic  theory.  Let  us  take,  for  example,  a  dilute 
solution  of  copper  sulphate,  into  which  hydrogen  sulphide  is  being 
passed.  Copper  sulphate  is  almost  entirely  ionized,  hydrogen 
sulphide  only  to  a  very  small  degree  (d).  We  therefore  have  in 
the  solution: 

Cu"+  S04"+2OT'  +  £S"  +  (1-£)H2S, 

the  cations  being  represented  by  a  point  and  the  anions  by  a  line 
above  and  to  the  right,  and  the  number  of  these  points  or  lines 
indicating  the  ionic  valence  (§  76). 

Some  of  the  copper  ions  and  sulphur  ions  will  then  unite  to 
form  undissociated  molecules,  CuS,  which  are  only  slightly  soluble 
in  water  and  are  therefore  precipitated.  As  S-ions  thus  disappear, 
the  equilibrium  between  hydrogen  sulphide  and  its  ions  is  dis- 
turbed; new  H2S  molecules  are  then  split  up  into  ions,  so  that 
there  are  again  S-ions  present,  which  can  unite  with  copper,  and 
so  on.  The  action  proceeds  according  to  the  equation: 

CuS04 + H2S = CuS  +  H2SO4, 

Insol. 

or,  if  only  the  ions  which  take  part  in  it  are  represented: 

Cu"+  S"  =  CuS. 

This  takes  place  quantitatively  if  the  copper  solution  is  dilute 
and  no  considerable  amount  of  any  strong  free  acid  was  added. 
However,  if  these  conditions  are  not  fulfilled  and,  as  a  result,  the 


73.]  HYDROGEN  SULPHIDE.  119 

concentration  of  the  hydrogen  ions  is  rather  high,  the  presence 
of  these  ions  reduces  the  ionization  of  H2S  so  much  (§  66)  that 
no  precipitate  can  be  formed.  "The  application  of  the  mass- 
action  law  to  the  case  is  very  simple.  Copper  sulphide,  when  in 
contact  with  water,  dissolves  to  an  extremely  small  extent;  in 
this  solution  we  have  the  equilibrium: 

Cu  '+  S"  *±  CuS. 

If  the  concentrations  of  the  two  ions  are  a  and  6,  and  that  of  the 
undissociated  -copper  sulphide  is  c}  we  have  the  equation 


k  being  a  constant  for  a  fixed  temperature  (§  66). 

The  product  ab  has  a  definite  value  for  every  saturated  solu- 
tion (since  c  is  definite).  This  value  is  known  as  the  solubility 
product  of  the  substance  in  question.  If  in  any  case  the  product 
ab  is  less  than  this  value,  none  of  the  substance  can  separate  out, 
because  the  solution  will  then  be  unsaturated;  if,  however,  the 
product  is  greater  than  the  solubility  product,  the  substance  will 
be  precipitated. 

As  soon  then  as  the  concentration  of  the  S-ions  becomes  so 
small  (because  of  the  reduction  of  the  ionization  of  hydrogen  sul- 
phide by  the  H-ions  of  the  acid)  that  it  makes  the  value  of  ab 
smaller  than  that  of  the  solubility  product  for  copper  sulphide,  no 
precipitate  will  be  formed.  If,  however,  the  liquid  is  diluted,  the 
concentration  of  the  H-ions  decreases;  then,  if  hydrogen  sulphide 
is  passed  in,  the  concentration  of  the  S-ions  increases.  The  value 
of  the  solubility  product  can  in  this  way  be  exceeded,  in  which 
event  copper  sulphide  will  be  precipitated. 

If  a  small  quantity  of  strong  acid  be  added  to  a  precipitate  of 
copper  sulphide  suspended  in  water,  only  a  very  small  amount 
of  the  sulphide  will  dissolve;  to  be  sure,  the  H-ions  of  the  strong 
acid  will  remove  a  part  of  the  S-ions,  yielding  some  undissociated 
hydrogen  sulphide,  so  that  in  order  to  establish  equilibrium  a  trace 
of  copper  sulphide  must  go  into  solution;  but  soon  the  point  will 
be  reached  when  so  many  Cu-  and  S-ions  are  again  in  the  solution 
that  the  value  of  the  solubility  product  is  reached.  After  this 
moment  no  more  copper  sulphide  goes  into  solution.  Since  the 
value  of  the  solubility  product  is  very  low,  the  solubility  of  the 


120  INORGANIC  CHEMISTRY.  [§§  73- 

sulphide  in  dilute  strong  acids  is  very  slight;    this  accounts  for 
the  practically  complete  precipitation  of  the  copper  sulphide. 

On  the  other  hand,  if  the  solubility  product  of  a  sulphide  is 
greater,  as  in  the  case  of  iron  sulphide,  the  addition  of  sulphuretted 
hydrogen  to  the  solution  of  an  iron  salt,  e.g.  ferrous  sulphate, 
FeSO4,  will  cause  no  precipitate  of  iron  sulphide,  and  iron  sulphide 
will,  unlike  the  previous  case,  be  dissolved  by  dilute  strong  acids. 
When  hydrogen  sulphide  is  led  into  a  solution  of  ferrous  sulphate 
to  the  point  of  saturation,  the  concentration  of  the  S-ions  is,  on 
account  of  the  slight  ionization  of  hydrogen  sulphide,  not  great 
enough  together  with  that  of  the  Fe-ions  to  reach  the  solubility 
product  of  iron  sulphide,  hence  no  precipitate  forms.  Moreover, 
when  dilute  hydrochloric  acid  is  added  to  iron  sulphide,  the  H-ions 
and  the  S-ions  form  undissociated  H2S  molecules  and  the  concen- 
tration of  the  S-ions  therefore  becomes  too  small  in  comparison 
with  the  value  of  the  solubility  product  to  prevent  solution ;  hence 
in  the  presence  of  enough  acid  all  the  iron  sulphide  goes  into  solution. 

It  now  becomes  clear,  too,  why  iron  solutions  are  precipitated 
by  ammonium  sulphide.  This  takes  place  according  to  the  equa- 
tion 

FeS04+  (NH4)2S  =  FeS  +  (NH4)2S04. 

In  this  case  there  are  no  H-ions  in  the  solution  to  act  on  the  iron 
sulphide. 

The  reason  for  the  non-precipitation  of  metals  like  barium, 
etc.,  either  by  sulphuretted  hydrogen  or  ammonium  sulphide  lies 
in  the  easy  solubility  of  their  sulphides. 

Hydrogen  Persulphide. 

74.  If  a  solution  of  sodium  sulphide,  Na-jS,  is  digested  with  sulphur, 
the  sulphur  dissolves  and  the  liquid  contains  compounds  called  poly- 
sulphides  and  having  formulae  from  Na2S2  up  to  Na,S5,  according  to  the 
amount  of  sulphur  employed.  On  pouring  such  a  solution  into  cold 
dilute  hydrochloric  acid  an  oil  separates  out  which,  by  distillation  under 
low  pressure,  yields  two  compounds,  having  the  formulae  H2S2  and  H2S3. 
The  hydrogen  disulphide  is  at  ordinary  temperatures  a  yellowish,  water- 
clear  liquid  with  a  consistency  somewhat  like  that  of  water.  With 
alkalies  it  decomposes  violently.  Under  ordinary  pressure  the  liquid 
boils  with  partial  stability  at  74-75°.  Its  specific  gravity  is  1.376.  Its 
fumes  attack  the  eyes  and  mucous  membranes  vigorously. 


75.]     COMPOUNDS  OF  SULPHUR  WITH  THE  HALOGENS.       121 

Hydrogen  trisulphide  at  ordinary  temperatures  is  a  bright  yellow 
liquid  somewhat  more  mobile  than  olive  oil.  Its  specific  gravity  at  15° 
is  1.496.  The  odor  reminds  one  of  sulphur  chloride  and  camphor.  The 
liquid  solidifies  at  —52°.  On  warming,  it  turns  darker,  becomes  more 
viscid  and,  at  about  90°,  begins  an  active  evolution  of  hydrogen  sulphide. 
Alkalies  produce  vigorous  decomposition. 

The  sensitiveness  of  these  compounds  toward  alkalies  is  so  great  that 
they  can  only  be  prepared  and  kept  in  glass  vessels  whose  inside  surfaces 
have  been  previously  freed  from  traces  of  alkali  by  treating  with  an  acid. 


Compounds  of  Sulphur  with  the  Halogens. 

75.  If  chlorine  is  conducted  over  molten  sulphur,  sulphur 
monochloride,  S2G12,  is  formed.  Its  formula  is  based  on  its  vapor 
density  and  analysis.  It  is  a  yellow  liquid  of  a  very  disagreeable, 
pungent  odor,  which  excites  one  to  tears.  It  boils  at  139°  and 
possesses  in  a  high  degree  the  ability  to  dissolve  sulphur — as  much 
as  66%  at  ordinary  temperatures.  This  solution  is  a  thick  syrupy 
liquid.  It  is  used  in  the  vulcanizing  of  rubber. 

Sulphur  monochloride  is  slowly  decomposed  by  water: 

2S2C12 + 2H2O  =  SO2  +  3S  +  4HC1. 

Two  other  compounds  of  sulphur  and  chlorine  are  known, 
SC12  and  SCU. 

Sulphur  dichloride,  SC12,  is  formed,  slowly,  when  sulphur 
monochloride  is  mixed  with  liquid  chlorine.  The  mixture  has  a 
yellow  color  at  first  but  after  a  few  days  it  turns  red.  A  determina- 
tion of  the  vapor  density  of  the  red  substance  and  of  its  lowering 
of  the  freezing-point  of  acetic  acid  or  benzene  leads  to  the  formula 
SC12.  It  should,  however,  be  borne  in  mind  that  a  mixture  of 
sulphur  monochloride  and  chlorine,  S2C12  +  C12,  must  give  the  same 
molecular  weight  as  the  compound  SC12.  The  existence  of  the 
SC12  compound  is  proved  not  only  by  the  above-mentioned  change 
of  color  of  mixtures  of  sulphur  monochloride  and  chlorine  but  also 
by  the  following  observations :  (1)  The  composition  of  the  vapor 
given  off  from  fresh  mixtures  of  sulphur  monochloride  and  chlorine 
is  entirely  different  from  that  of  the  vapor  given  off  after  the  mix- 
ture has  turned  red.  (2)  Mixtures  of  sulphur  monochloride  and 


122       SB  INORGANIC  CHEMISTRY.  [§§75- 


chlorine  decrease  in  volume,  and  this  diminution  is  greatest  when 
the  composition  corresponds  to  S2C12  +  C12.  (3)  On  distilling 
under  4  mm.  pressure  80-90%  of  the  liquid  could  be  distilled  over 
almost  constant  at  —24°.  At  ordinary  pressure  the  compound 
boils  at  about  59-60°  with  decomposition. 

Sulphur  tetrachloride,  SC14,  can  be  obtained  as  a  fine  white 
powder,  apparently  not  crystalline,  when  a  chlorine-sulphur 
mixture  of  the  composition  S  +  CU  or  S  +  Cle  is  cooled  to  —75° 
(where  it  solidifies)  and  then  centrifuged.  The  tetrachloride  melts 
at  about  —33°,  giving  off  chlorine  abundantly.  The  exact  tem- 
perature could  not  be  determined.  Here,  too,  the  form  of  the 
freezing-point  curve  indicates  unquestionably  the  formula  of  the 
substance,  viz.,  SC14.  Cf.  §  237. 

With  bromine  and  iodine  sulphur  gives  analogous  com- 
pounds. 

Fluorine  unites  with  sulphur  to  form  a  gas  of  the  formula  SFe 
and  of  rather  surprising  properties.  It  is  colorless,  odorless  and 
incombustible.  At  —  55°  it  solidifies  with  the  formation  of  crystals. 
Notwithstanding  its  high  percentage  of  fluorine  it  is  chemically 
so  indifferent  that  it  almost  resembles  nitrogen  in  this  respect 
(see  p.  164).  For  instance,  it  is  not  decomposed  by  fused  alkalies 
nor  by  copper  oxide  at  dull-red  heat.  It  can  be  heated  with 
hydrogen  without  yielding  hydrogen  fluoride.  Moreover,  sodium 
can  be  fused  in  sulphur  hexafluoride  without  losing  its  metallic 
surface,  the  gas  not  being  attacked  by  the  metal  till  the  boiling- 
point  of  the  latter  is  reached. 

VALENCE. 

76.  Certain  elements  have  the  property  whereby  their  atoms 
can  combine  with  only  one  atom  of  another  element.  The  halogens 
on  the  one  hand  and  hydrogen  on  the  other  are  able  to  form  only 
compounds  of  the  type  HX(X  =  halogen).  This  property  of  the 
atoms  is  called  univalence. 

In  the  case  of  other  elements  like  oxygen  and  sulphur  each 
atom  can  enter  into  compounds  with  two  univalent  atoms  (exam- 
ples: H2S,  H20).  These  are  therefore  called  bivalent. 

The  number  of  univalent  atoms  that  can  combine  with  one 
atom  of  a  given  element  serves  in  an  analogous  way  as  a  measure 
of  valence  in  general.  An  atom  of  nitrogen,  for  instance,  unites 


76.]  VALENCE.  123 

with  three  atoms  of  hydrogen;    nitrogen  is  therefore  trivalent; 
carbon  is  quadrivalent,  etc. 

/H 
The  valence  is  ordinarily  indicated  by  lines,  as  in  O\       and 

/H  XH 

N — H,  each  line  representing  a  valence  unit  (unit  bond). 
\H 

The  valence  of  one  and  the  same  element  may  be  different 
according  to  the  nature  of  the  univalent  elements  with  which  it 
combines.  Sulphur,  for  instance,  can  only  unite  with  two  hydrogen 
atoms,  but  with  univalent  chlorine  it  forms  the  compound  SC14, 
with  fluorine  even  SF6.  The  valence  of  sulphur  in  these  cases  is 
therefore  four  and  six.  The  preparation  of  sulphur  compounds 
with  more  than  six  univalent  atoms  has  not  yet  been  accomplished ; 
hence  its  maximum  valence  is  six. 

The  halogens  are  univalent  towards  hydrogen,  but  in  relation 
to  each  other  they  display  more  than  one  valence,  as  may  be  seen 
from  the  compounds  IC13  and  IC15 ;  in  the  compound  C^Oy  (§  60) 
the  maximum  valence  of  chlorine  can  even  be  assumed  to  be  seven. 

It  has  been  very  generally  observed  that  when  the  maximum 
valence  of  an  element  is  an  even  or  an  uneven  number,  its  lower 
valences  are  of  the  same  sort;  the  halogens  and  sulphur  illustrate 
this.  However,  these  are  exceptions  to  this  rule. 

The  valence  also  depends  upon  the  temperature.  We  shall  soon  see  that 
SO3  dissociates  at  a  high  temperature  into  SO2  and  oxygen;  while  sulphur  is 

f  //\ 

sexivalent  towards  oxygen  at  lower  temperatures  (  S=O  1,  it'  is  only  quadri- 

\  \o/ 

/    S°\ 
valent  towards  oxygen  above  700°  fSr'     J .     The  valence  must  also  depend 

on  the  pressure,  for  the  latter  exerts  a  great  influence  on  the  dissociation. 

The  basis  for  the  above  sort 'of  formulae  is  the  idea,  borrowed  from  organic 
chemistry,  that  the  atoms  of  a  molecule  may  not  assume  any  conceivable 
arrangement  whatsoever,  but  that  there  is  a  definite  order  in  every  molecule. 

For  some  extensions  of  the  idea  of  valence  see  §  317. 

Valence  of  Ions. — In  the  solution  of  an  electrolyte  the  sums  of 
all  the  positive  and  all  the  negative  amounts  of  electricity  must 
be  equal,  for  the  solution  acts  as  electrically  neutral.  In  a  solu- 
tion of  hydrochloric  acid  the  positive  charge  of  the  H-ions  must 


124  INORGANIC    CHEMISTRY.  [§§76- 

be  numerically  equal  to  the  negative  charge  of  the  Cl-ions  and, 
since  the  same  number  of  both  ions  are  present,  each  Cl-ion  must 
carry  a  charge  equal,  but  opposite  in  sign,  to  that  of  an  H-ion. 
In  a  sulphuric  acid  solution,  however,  the  two  H-ions  together 
must  possess  just  as  much  positive  electricity  as  the  SO4-ion  nega- 
tive electricity.  The  SO/'  ion  is  therefore  called  bivalent  in  re- 
spect to  the  hydrogen  ion.  It  is  readily  seen  how  the  valence  of 
other  ions  can  be  determined  in  an  analogous  manner,  for  it  is 
equal  to  the  numerical  value  of  their  electrical  charge,  that  of 
the  hydrogen  ion  being  taken  as  unity. 

Compounds  of  Sulphur  with  Oxygen. 

77.  Of  those  containing  only  the  two  elements  three  are  known, 
viz.,  S20a,  SO 2,  and  SO3.  Especial  importance  attaches  itself, 
however,  only  to  SO2  and  SOs;  the  two  others  have  been  little 
studied. 

Sulphur  Sesquioxide,  S2O3. 

This  is  obtained  when  sulphur  is  treated  with  its  trioxide.  It  is  a 
blue  liquid,  which  congeals  to  a  malachite-green  mass  and  is  soluble  in 
fuming  sulphuric  acid,  giving  a  blue  solution.  On  being  warmed  it 
breaks  up  into  sulphur  and  the  dioxide : 

2S203=3S02  +  S. 

Water  decomposes  it  with  the  formation  of  sulphur,  sulphurous  acid  and 
polythionic  acids. 


SULPHUR  DIOXIDE,  SULPHUROUS  ANHYDRIDE,  S02. 

78.  This  gas  occurs  in  nature  in  volcanic  gases.  It  is  formed 
when  sulphur  burns  in  the  air  or  in  oxygen;  the  well-known  odor 
of  burning  sulphur  is  due  to  it.  A  little  trioxide  is  also  formed 
by  this  combustion.  The  laboratory  method  of  preparation  con- 
sists in  decomposing  sulphuric  acid  with  copper. 

2H2S04 + Cu  =  CuSO4 + SO2  +  2H20. 

For  this  purpose  concentrated  sulphuric  acid  is  heated  with  cop- 
per turnings,  no  action  taking  place  at  ordinary  temperatures. 
The  process  can  be  explained  by  supposing  that  at  the  high  tern- 


78.]  SULPHUR  DIOXIDE.  125 

perature  of  the  reaction  copper  is  oxidized  by  sulphuric  acid  to 
copper  oxide  with  the  evolution  of  sulphur  dioxide: 


The  copper  oxide  reacts  of  course  with  a  second  molecule  of 
sulphuric  acid,  producing  copper  sulphate. 

•  The  reduction  of  concentrated  sulphuric  acid  by  heating  with 
charcoal  is  also  a  convenient  method  of  preparation: 

2H2SO4  +  C  =  2H20  +  2S02  +  C02. 

However,  as  this  equation  shows,  the  gas  is  obtained  mixed  with 
one  third  of  its  volume  of  carbon  dioxide,  from  which  it  cannot  be 
separated  directly. 

Moreover,  sulphur  dioxide  can  be  obtained  by  the  action  of 
oxygen  on  sulphur  compounds,  thus,  e.g.  by  the  roasting  of  pyrite 
in  a  current  of  air: 

FeS2  +  302  =  SO2  +  FeS04. 

Pyrite. 

This  reaction  is  employed  on  a  large  scale  in  the  commercial 
manufacture  of  sulphuric  acid. 

The  action  of  sulphur  on  oxygen  compounds  also  yields  sul- 
phurous oxide,  e.g.  heating  copper  oxide  or  manganese  dioxide 
with  sulphur: 

2CuO  +  2S  =  Cu2S  +  S02  ;    Mn02  +  2S  =  MnS  +  S02. 

Finally,  the  dioxide  is  also  formed  by  heating  an  oxygen  com- 
pound (CuO)  with  a  sulphur  compound  (CuS)  : 

CuS+2CuO  =  3Cu+S02. 

Physical  Properties.  —  At  ordinary  temperatures  and  pressures 
sulphur  dioxide  is  a  gas.  It  has  a  peculiar  taste  and  odor.  It  is 
easily  liquefied,  the  boiling-point  being  —8°.  Its  evaporation 
produces  a  marked  depression  of  temperature,  sometimes  extend- 
ing to  —50°;  at  —76°  it  becomes  solid.  Liquid  sulphur  dioxide 
dissolves  many  salts,  in  some  cases  with  a  characteristic  color.  It 


126  INORGANIC  CHEMISTRY.  [§  78- 

is  very  soluble  in  water;  at  0°  1  vol.  H2O  dissolves  79.79  vols. 
S02,  at  20°  39.37  vols.  S02.  Boiling  the  solution  expels  all  the 
gas  (§  83). 

Chemical  Properties. — Sulphur  dioxide  is  an  acid  anhydride; 
its  aqueous  solution  has  an  acid  reaction  and  behaves  in  general 
like  that  of  an  acid  (§  83).  It  is  easily  oxidized  by  oxidizing- 
agents  to  the  trioxide.  This  occurs,  for  instance,  when  a  mixture 
of  sulphur  dioxide  and  air  or  oxygen  is  passed  over  hot  spongy 
platinum  or  platinum  asbestos.  In  aqueous  solution  this  oxidation 
takes  place  readily  at  ordinary  temperatures.  The  oxidation  of 
the  dioxide  can  also  be  brought  about  by  chlorine-water,  bromine 
and  iodine : 

C12+2H20  +  S02  =H2S04+2HC1; 

also  by  chromic  acid,  which  is  reduced  to  chromium  sulphate,  or 
by  potassium  permanganate,  which  is  reduced  to  a  mixture  of 
manganese  and  potassium  sulphates,  and  therefore  loses  its  color: 

2KMnO4 + 5S02 + 2H2O  =  K2SO4 + 2MnSO4 + 2H2S04. 

Lead  peroxide  glows  faintly  in  a  current  of  sulphur  dioxide 
and  is  reduced  to  lead  sulphate — from  brown  to  white: 

Pb02  +  S02=PbS04. 

It  is  to  its  reducing  action  that  the  bleaching  effect  of  sul- 
phurous oxide  on  vegetable  coloring-matters  is  due.  A  red  rose, 
for  example,  loses  color  in  it.  The  gas  probably  reacts  with 
water,  setting  hydrogen  free,  which  latter  effects  the  reduction 
and  hence  the  bleaching: 

S02+2H2O  =H2S04+H2. 

In  this  case,  therefore,  bleaching  depends  on  a  reduction;  as  a 
matter  of  fact  the  color  returns  in  many  instances,  when  the 
bleached  article  is  exposed  to  the  oxidizing  action  of  the  air.  Silk, 
wool  and  straw,  i.e.  substances  that  cannot  stand  the  chlorine 
bleaching,  are  whitened  commercially  with  sulphurous  oxide. 
It  also  finds  use  as  an  antiseptic. 


78.]  SULPHUR   DIOXIDE.  127 

The  reduction  of  iodic  acid  by  sulphur  dioxide  is  sometimes  employed 
as  a  test  for  the  latter.  For  this  purpose  strips  of  paper  are  dipped  in  a 
solution  of  potassium  iodate  and  starch,  which  turns  blue  in  the  presence 
of  sulphur  dioxide — iodine  being  set  free  (§  47). 

If  the  reaction  is  carried  out  in  dilute  solution,  a  peculiar  phenomenon 
is  observed ;  the  blue  color  of  starch  iodide  does  not  appear  directly  when 
the  solutions  of  sulphur  dioxide  and  iodic  acid  are  mixed,  but  is  with- 
held for  a  certain  number  of  seconds  (definite  for  every  concentration 
at  constant  temperature),  when  it  suddenly  appears.  The  following 
reactions  come  into  play : 

I.  3SO2aq  +  HIO3  =  3H2SO4aq  +  HI. 

The  hydriodic  acid  thus  formed  is  at  once  oxidized  by  the  iodic  acid 
still  present : 

II.  5HI  +  HI03=3H2O  +  6I. 

So  long  as  sulphur  dioxide  is  present,  it  reduces  the  iodine  in  this 
dilute  solution  to  hydriodic  acid: 

III.  21  +  SO2aq  +  2H2O  =  H2SO4aq + 2HI. 

Not  until  all  the  dioxide  is  used  up  by  the  reactions  I  and  III  does  the 
free  iodine  suddenly  appear  according  to  II. 

There  are  some  substances  which  are  able  to  extract  oxygen 
from  sulphur  dioxide,  i.e.  the  latter  can  also  act  as  an  oxidizing- 
agent.  Ignited  magnesium  ribbon  continues  to  burn  in  sulphur 
dioxide,  forming  magnesium  oxide  and  sulphur.  Hydrogen  sul- 
phide and  sulphur  dioxide  have  respectively  an  oxidizing  and  a 
reducing  effect  on  each  other,  which  follows  mainly  the  equation: 

2H2S  +  SO2=2H2O  +  3S. 

Sulphur  dioxide  is  decomposed  by  electric  sparks  into  sulphur 
and  the  trioxide. 

The  action  of  the  electric  sparks  is  to  be  ascribed  solely  to  the  sudden 
and  enormous  rise  of  temperature  which  they  produce  and  the  rapid  cool- 
ing that  immediately  follows,  for  the  gas  particles  which  have  become 
heated  by  the  sparks  are  immediately  cooled  again  by  surrounding 
objects.  As  a  result  the  products  formed  do  not  have  time  to  react  in 
the  opposite  direction.  The  correctness  of  this  view  was  demonstrated 
by  ST.  CLAIRE  DEVILLE  with  the  help  of  an  apparatus  which  made  it  pos- 


128  INORGANIC    CHEMISTRY. 


[§§78- 


sible  to  cool  objects  very  rapidly  from  a  very  high  temperature.  This 
apparatus,  the  cold-hot  tube,  consists  of  a  rather  wide  porcelain 
tube,  which  is  heated  to  a  bright  glow  in  a  furnace  and  which  contains 
a  concentric  thinner  metallic  tube,  through  which  cold  water  is  forced  so 
rapidly  that  the  tube  maintains  a  low  temperature.  When  DEVILLE 
introduced  sulphur  dioxide  into  the  space  between  the  two  tubes,  it  was 
seen  after  some  time  that  the  inner  tube,  which  was  made  of  silver-plated 
copper,  had  turned  black  because  of  the  formation  of  silver  sulphide, 
while  at  the  same  time  the  formation  of  sulphur  trioxide  could  be  detected 
(by  its  producing  sulphuric  acid  with  water,  a  precipitate  being  given  by 
barium  chloride). 

The  composition  of  sulphurous  oxide  can  be  determined  in  the 
following  manner:  When  sulphur  burns  in  oxygen  no  change  of 
volume  is  observed  after  cooling.  Therefore  just  as  many  mole- 
cules of  sulphurous  oxide  have  been  formed  as  oxygen  molecules 
consumed.  The  sulphurous  oxide  molecule  must  therefore  con- 
tain two  atoms  of  oxygen.  The  specific  gravity  of  the  gas  has 
been  found  to  be  2.2639  (air  =  l),  or  32.6  (O  =  16),  so  that  its 
molecular  weight  is  65.2.  If  we  subtract  2X16  from  this  for 
two  atoms  of  oxygen,  there  remains  33.2  for  sulphur,  the  atomic 
weight  of  which  is  32.  We  thus  see  that  only  one  atom  of  sul- 
phur is  present  in  the  molecule  of  sulphurous  oxide  and  that  the 
formula  of  the  latter  is  862. 


SULPHUR   TRIOXIDE,  SULPHURIC  ANHYDRIDE,  S03. 

79.  This  compound  is  found  in  a  small  amount  in  the  fumes 
of  burning  sulphur  (§  78).  As  was  stated  above,  oxygen  and 
sulphur  dioxide  unite  to  form  the  trioxide  in  the  presence  of 
platinized  asbestos.  On  the  other  hand,  the  trioxide  breaks  up 
into  the  dioxide  and  oxygen  at  an  elevated  temperature,  so  that 
the  formation  of  the  trioxide  from  the  dioxide  and  oxygen  is 
evidently  a  reversible  process,  which  is  expressed  thus: 


If  we  call  the  pressure  of  SO2  in  the  equilibrium  condition  pi,  that 


79.  SULPHUR  TRIOXIDE.  129 

of  O2  p2  and  that  of  SO3  p3,  it  follows  from  §51  that  the  equilibrium 
relation  is  expressed  by 


where  K  is  the  equilibrium  constant. 

The  combination  of  sulphur  dioxide  and  oxygen  is  easily 
accomplished  (in  the  presence  of  platinum)  at  about  500°;  that 
is  to  say,  the  above  equilibrium  is  shifted  almost  wholly  to  the 
right  at  this  temperature.  If  the  temperature  is  raised,  the 
dissociation  of  trioxide  begins  and  at  about  1000°  it  is  com- 
plete. 

The  union  of  S02  and  O2  also  occurs  under  the  influence  of  ultra- 
violet rays.  These  rays  are  best  produced  by  a  quartz-mercury  arc  lamp. 
The  gases  that  are  to  be  exposed  to  the  rays  must  also  be  contained  in 
quartz  vessels,  since  glass  is  opaque  to  ultraviolet  rays. 

Furthermore,  an  equilibrium  2SO2+O2<=±2SO3  also  establishes  itself 
under  the  action  of  these  rays;  for  not  only  is  the  union  of  SO2  and  02 
incomplete,  but  SO3,  on  the  other  hand,  breaks  up  under  the  same 
experimental  conditions,  yielding  the  same  equilibrium  mixture.  This 
light  equilibrium  differs,  however,  in  many  respects  from  the  "  tem- 
perature equilibrium."  In  the  first  place,  S03,  in  the  presence  of 
platinum,  does  not  begin  to  dissociate  perceptibly  until  300°.  The  light 
influence,  however,  is  evident  even  at  room  temperature.  Like  the 
effect  of  catalyzers,  the  action  of  light  is  retarded  by  sufficiently  careful 
drying  of  the  gases.  There  is  an  optimum  moisture  content  for  the  con- 
tact process,  but  the  action  of  light  is  effective  even  when  the  gases  are 
passed  through  the  illumination  vessel  in  a  very  moist  condition.  The 
light  equilibrium  is  not  perceptibly  affected  by  a  marked  change  of  tem- 
perature, but  it  is  very  sensitive  to  varying  intensity  of  illumination. 
Just  exactly  as  the  dissociation  increases  with  rising  temperature,  so  it 
increases  as  the  illumination  grows  stronger.  The  reader  can  form  an 
idea  of  the  extent  of  the  decomposition  from  the  observation  of  COEHN 
and  BECKER  that,  with  a  mercury  lamp  consuming  9  amp.,  the  equi- 
librium established  itself  when  about  35%  of  the  SO3  was  decom- 


Sulphur  trioxide  can  also  be  obtained  by  heating  certain  sul- 
phates; in  the  arts  ferric  sulphate  is  thus  used: 


130  INORGANIC  CHEMISTRY.  [§§80- 

"  Fuming  sulphuric  acid"  (oleum)  is  a  solution  of 
sulphur  trioxide  in  sulphuric  acid ;  the  anhydride  can  be  obtained 
from  it  by  distillation. 

80.  Physical  Properties. — Perfectly  dry  sulphur  trioxide  melts  at 
17.7°  and  boils  at  46°.  It  looks  much  like  ice  but  usually  appears 
in  another  modification,  viz.,  long  asbestos-like  needles  with  a 
silky  lustre.  These  crystals  have  no  sharp  melting-point  but 
sublime  on  heating.  This  modification  is  the  stable  one,  for  the 
other  goes  over  into  it  spontaneously.  This  transformation  is 
greatly  accelerated  by  traces  of  water.  The  asbestine  modifica- 
tion consists  of  double  molecules  (863)2,  the  glacial  form  of  simple 
molecules  (863).  This  is  shown  by  the  depression  of  the  freezing- 
point  of  phosphorus  oxychloride.  The  first  is  therefore  called  a 
polymer  of  the  second.  It  is  also  worth  noting  that  the  863 
modification  is  very  readily  soluble  in  concentrated  sulphuric 
acid,  while  the  other,  (863)2,  dissolves  with  difficulty. 

Chemical  Properties. — Sulphur  trioxide  unites  very  easily  with 
water  to  form  sulphuric  acid: 

SO3+H20=H2S04. 

It  therefore  fumes  vigorously  when  exposed  to  moist  air.  On 
introducing  it  into  water,  combination  and  great  evolution  of  heat, 
accompanied  by  sizzling,  results.  It  reacts  energetically  with 
many  metallic  oxides  also,  forming  sulphates.  Baryta,  for  exam- 
ple, glows  in  contact  with  it.  When  its  vapor  is  passed  through 
a  red-hot  tube,  it  is  decomposed  into  the  dioxide  and  oxygen. 

Composition. — The  decomposition  just  mentioned  permits  us 
to  establish  the  composition  of  sulphuric  oxide.  The  dissociation 
products,  862  and  62,  are  formed  in  the  volume  ratio  2:1.  Now 
the  specific  gravity  of  sulphuric  oxide  is  2.75  (air  =  l),  from  which 
the  molecular  weight  is  calculated  to  be  79.1.  This  figure  corre- 
sponds to  the  formula  863  (32  +  3X16)  and  it  also  harmonizes 
with  the  above  dissociation;  for  it  is  clear  that  2  vols.  863  must 
then  yield  2  vols.  862  and  1  vol.  02 : 

2S03=2SO2  +  02. 

2  vols.       2  vols.     1  vol. 


82.]  OXYGEN    ACIDS  OF  SULPHUR.  131 

Oxygen  Acids  of  Sulphur. 

81.  Sulphur  forms  an  unusually  large  number  of  acids  with 
oxygen  and  hydrogen,  namely  nine.  They  are  as  follows: 

1.  Thiosulphuric  acid H2S203. 

2.  Hyposulphurous  acid H2S204. 

3.  Sulphurous  acid H2SO3. 

4.  Sulphuric  acid H2S04. 

5.  Persulphuric  acid H2S208. 

6.  Dithionic  acid H2S206. 

7.  Trithionic  acid H2S306. 

8.  Tetrathionic  acid H2S406. 

9.  Pentathionic  acid H2S506. 

It  is  an  important  fact,  however,  that  of  these  nine  acids  only 
sulphuric  acid  has  really  been  isolated;  all  the  others  are  known 
only  hi  aqueous  solution  or  hi  the  form  of  salts.  The  two  hydro- 
gen atoms  which  each  of  these  acids  possesses  are  both  replaceable 
by  metals;  they  are  therefore  dibasic  acids.  With  such  acids  it 
is  possible  that  just  one  of  the  hydrogen  atoms  be  replaced  by  a 
metal.  The  salts  thus  formed  are  called  acid  salts. 

By  different  methods,  e.g.  the  cryoscopic  method,  it  is  found 
that  the  aqueous  solution  of  dibasic  acids  AH2  contains  chiefly  the 
ions  H"  and  HA';  it  is  only  when  these  solutions  are  very  dilute 
that  the  anion  HA'  splits  up  further  into  H*  and  A".  In  the  case 
of  the  M2A  salts,  however,  there  is  an  ionization  into  2M*+ A"; 
but  in  that  of  the  acid  salts  MHA  the  ions  are  chiefly  M*  and  HA'. 
How  far  the  anion  HA'  is  split  up  does  not  depend  merely  on  the 
concentration,  but  also  to  a  considerable  degree  on  the  strength 
of  the  acid,  HA'  being  more  ionized  in  strong  than  in  weak  acids 
of  the  same  concentration. 


THIOSULPHURIC  ACID,  H2S203. 

82.  This  acid  can  only  exist  in  dilute  aqueous  solution  and  is 
even  then  very  unstable,  decomposing  completely  in  a  short  time. 
The  salts  are,  however,  stable  and  can  be  prepared  in  the  following 
ways: 


132  INORGANIC  CHEMISTRY.  [§§82- 

1.  By  boiling  the  solution  of  a  sulphite  with  sulphur; 


Sodium  sulphite. 

or  S03"  +  8  =  8203'', 

only  the  anion  being  changed. 

2.  By  the  oxidation  of  sulphides  in  the  air* 


2=2CaS203. 

Calcium  disulphide. 

3.  By  the  action  of  sulphur  dioxide  on  the  solution  of  a  sul- 
phide: 

4Na2S  +  6SO2  =  4Na2S2O3  +  S2. 

The  most  important  salt  is  the  sodium  thiosulphate,  formerly 
and  even  yet  often  called  sodium  hyposulphite,  or,  abbreviated, 
"hypo."  It  is  very  soluble  in  water;  the  solution,  when  used  in 
excess,  has  the  property  of  dissolving  readily  the  halogen  com- 
pounds of  silver,  hence  its  extensive  use  in  photography  (§  247). 
It  is  easily  oxidized  by  oxidizing-agents,  usually  to  the  sulphate. 
This  takes  place  with  potassium  permanganate,  nitric  acid  and 
chlorine,  for  example.  Practical  use  is  also  made  of  this  latter 
property  by  employing  sodium  thiosulphate  as  an  antichlor  in 
bleaching,  i.e.  to  remove  the  last  traces  of  chlorine  which  cling  to 
the  bleached  material  very  obstinately  and  have  an  injurious  effect. 

When  a  dilute  acid  is  added  to  a  dilute  solution  of  sodium 
thiosulphate,  the  following  decomposition  takes  place: 

=  2NaCl  +  H20  +  SO2  +  S  ; 


or  S203"+  2H'  =  HS03'+  H'  +  S. 

iipi 

id. 


Anion  of  sulphur- 
ous acic 


It  may  be,  however,  that  the  ions  first  unite  partially  to  form 
H2S203,  which  splits  up  into  H2O,  S  and  SO2. 

It  is  an  interesting  fact  that  in  this  decomposition  in  a  dilute  solution 
the  sulphur  precipitate  is  not  at  once  visible,  being  first  noticeable  after 
some  seconds,  or  even  minutes,  according  to  the  dilution.  It  was  formerly 
supposed  that  the  thiosulphu'ric  acid  remained  entirely  unchanged  until 
the  appearance  of  the  sulphur  and  the  decomposition  first  began  at  this 


84.]  HYPOSULPHUROUS,  AXD  SULPHUROUS   ACIDS.          133 

moment.  This  is, '  however,  incorrect;  for  when  a  dilute  solution  of 
thiosulphate  is  treated  with  an  equivalent  amount  of  dilute  acid  and 
the  solution  again  neutralized  before  the  appearance  of  the  sulphur 
deposit,  it  is  found  that  the  latter  appears  nevertheless  after  some  time. 
A  certain  part  of  the  free  thiosulphuric  acid  must  therefore  have  already 
decomposed,  but  the  sulphur  was  in  a  so  very  finely  divided  atate  in  the 
liquid  that  it  could  not  at  once  be  detected, — not  until  it  !aad  gathered 
together  to  form  larger  particles. 


Hyposulphurous  Acid, 

83.  As  early  as  the  18th  century  it  was  observed  that  zinc  is  dis- 
solved by  a  solution  of  sulphur  dioxide  in  water  without  the  evolution  of 
hydrogen.  SCHUTZENBERGER  was,  however,  the  first  to  show  that  a 
particular  acid  is  formed  thereby.  A  salt  of  this  acid  is  produced  by  the 
action  of  zinc  on  a  solution  of  acid  sodium  sulphite,  NaHS03,  or  by  the 
electrolysis  of  such  a  solution,  the  nascent  hydrogen  acting  as  a  reducing- 
agent.. 

Hyposulphurous  acid,  as  well  as  its  salts,  is  characterized  by  a  vigorous 
reducing  power.  It  precipitates  the  metals  from  solutions  of  sublimate 
(HgCl2),  silver  nitrate  and  copper  sulphate.  Iodine  solution  is  bleached 
by  it  with  the  formation  of  hydrogen  iodide  ;  indigo  is  reduced  to  indigo- 
white.  The  solution  is  also  very  easily  oxidized  by  free  oxygen.  It  is 
therefore  used  to  determine  the  amount  of  oxygen  dissolved  in  water. 
For  this  reason  it  must  be  kept  in  well-stoppered  vessels. 

BERNTHSEN  succeeded  in  preparing  the  solid  sodium  salt,  which 
proved  to  have  the  composition  Na2S204  +  2H20,  so  that  the  acid  itself 
has  the  formula  H2S2O4.  This  salt  was  isolated  by  preparing  a  concen- 
trated solution  of  it  and  precipitating  it  by  the  addition  of  a  suitable 
amount  of  solid  common  salt.  The  above  formula  is  also  confirmed  by 
a  direct  synthesis  of  the  sodium  salt  by  MOISSAN,  who  obtained  it  by  the 
action  of  dry  sulphur  dioxide  on  sodium: 


SULPHUROUS  ACID,  H2S03. 

84.  It  is  taken  for  granted  that  the  aqueous  solution  of  sulphur 
dioxide  contains  sulphurous  acid,  H2SO3,  for  this  solution  reacts 
acid,  conducts  the  electric  current,  gives  salts  with  bases  and 
evolves  hydrogen  with  some  metals,  e.g.  magnesium.  The  solution 
of  sulphur  dioxide  in  water  does  not  conform  to  the  law  of  HENRY 
(§  9)  at  ordinary  temperatures,  which  proves  that  a  combina- 
tion with  the  solvent  has  taken  place.  At  higher  temperatures, 


134  INORGANIC   CHEMISTRY.  [§§84- 

however,  the  solution  obeys  this  law  pretty  well.  A  fact  in  con- 
firmation of  this  is  that  all  the  sulphur  dioxide  can  be  expelled  from 
the  solution  by  boiling  it,  the  combination  being  then  wholly 
destroyed.  The  compound  H2SO3  itself  has,  however,  not  yet  been 
isolated.  The  salts  have  the  composition  M2SO3  and  MHSOs 
(M  being  an  atom  of  a  univalent  metal) .  The  acid  salts  are  almost 
all  soluble  in  water,  while  of  the  neutral  salts  only  those  of  the 
alkalies  are  soluble.  The  acid  sodium  sulphite,  NaHSO3  (sodium 
bisulphite),  is  frequently  employed  in  organic  chemistry.  Sul- 
phites in  solution  gradually  absorb  oxygen  from  the  air,  form- 
ing sulphates.  It  is  a  very  strange  fact  that  minute  quantities  of 
organic  substances,  e.g.  only  0.1%  of  alcohol  and  as  little  as  10~5 
gram  molecule  of  stannous  chloride,  greatly  hinder  this  oxida- 
tion. We  have  here  one  of  the  few  examples  of  a  retarding 
catalytical  action.  On  the  other  hand,  traces  of  copper  sul- 
phate considerably  accelerate  the  oxidation. 


SULPHURIC    ACID,  H2S04. 

85.  Sulphuric  acid  is  the  most  important  acid  of  sulphur.     It 
can  be  obtained  in  various   ways;  in  the   first   place  by  direct 
synthesis  from  its  elements.     According  to  §  79  sulphur  trioxide 
can  be  formed  directly  from  sulphur  and  oxygen,  and  this  yields 
sulphuric  acid  on  the  addition  of  water. 

The  acid  can  be  obtained  from  its  salts  by  distilling  them  with 
phosphoric  acid.  Its  formation  from  the  action  of  oxygen  on 
sulphur  compounds  is  illustrated  by  the  oxidation  of  an  aqueous 
SCVsolution  by  the  air.  On  the  other  hand  the  action  of  sulphur 
on  oxygen  compounds  may  also  give  sulphuric  acid;  thus  it  is 
formed  when  concentrated  nitric  acid,  HNOs,  is  boiled  with  sul- 
phur; and  again,  potassium  sulphate  is  formed  by  heating  sulphur 
with  saltpetre  (KNOs). 

86.  For  the    commercial   manufacture   of   sulphuric    acid    two 
processes  are  now  in  use,  the  lead-chamber  process  and  the  con- 
tact process.     Enormous  amounts  of  the  acid  are  produced  by 
these  two  methods. 

The  lead-chamber  process  is  based  on  the  follow- 
ing reactions:  1.  the  oxidation  of  sulphur  dioxide  by  nitric  acid 
in  the  presence  of  water;  2.  the  oxidation  by  the  oxygen  in  the 


86.]  SULPHURIC  ACID.  135 

air  of  lower  oxides  of  nitrogen  formed  from  the  nitric  acid  in  the 
previous  reaction.  These  are  partly  reconverted  to  nitric  acid 
and  partly  changed  to  certain  stages  of  oxidation  of  nitrogen 
which  oxidize  sulphur  dioxide  anew  to  sulphuric  acid.  By  this 
last  process  the  lower  nitrogen  oxides  are  again  formed,  but  are 
soon  reoxidized  by  atmospheric  oxygen  and  so  on.  One  might 
suppose  that  a  certain  amount  of  nitric  acid  would  suffice  to  con- 
vert unlimited  amounts  of  sulphur  dioxide  into  sulphuric  acid 
with  the  aid  of  the  air.  In  practice  this  is  not  true,  however;  for 
the  nitrogen  oxides  are  to  a  small  extent  still  farther  reduced  by 
sulphurous  oxide,  so  that  nitrous  oxide  or  nitrogen  are  formed, 
and  these  are  no  longer  able,  under  the  conditions  of  the  indus- 
trial process,  to  combine  with  oxygen. 

The  chemical  processes  which  lie  at  the  basis  of  the  manu- 
facture of  sulphuric  acid  will  be  taken  up  a  little  later  (§  128). 

From  a  technical  standpoint  the  lead-chamber  process  falls 
into  three  separate  parts: 

1.  The  preparation  of  sulphur  dioxide; 

2.  The  oxidation  of  sulphur  dioxide; 

3.  The  concentration  of  the  resulting  acid. 

(1)  The  material  for  the  production  of  the  dioxide  is  sulphur 
or  pyrite  (iron  pyrites,  FeS2).     Sulphur  yields  a  purer  acid  than 
pyrite;    that   prepared  from  the  latter  almost  always  contains 
arsenic.    The  roasting  of  the  pyrite  is  carried  on  in  furnaces,  the 
construction  of  which  varies  considerably.     In  all  of  them,  how- 
ever, the  sulphur  dioxide  leaves  the  furnace  mixed  with  a  good 
deal  of  air.     The  furnace  gases  pass  through  a  canal  in  which  the 
dust  particles  carried  along  by  the  draught  are  deposited. 

(2)  The  oxidation  of  the  sulphurous  acid  is  carried  out  in  ? 
structure   consisting  chiefly  of  three    parts,   the   Glover  tower, 
the   lead    chambers,    and    the    Gay-Lussac    tower.      The    gases 
enter    the    bottom    of   the    Glover   Tower,    which    is   made    of 
sheet  lead  lined  with  acid-proof  brick.     It  is  filled   with  lump 
stone,  over  which  is  laid  a  layer  of  smaller  pieces  of  coke.     On 
top  of  the  tower  is  a  reservoir  for  collecting  the  nitroso  sulphuric 
acid  (see  below)  that  conies  from  the  Gay-Lussac  tower  and  the 
lead  chambers  and  is  to  be  concentrated  in  the  Glover  tower.     It 
flows  down  over  the  stone  in  the  tower  from  a  horizontally  revolv- 
ing tube.     From  the  Glover  tower  the  gases  enter  the  lead  cham- 


136  INORGANIC  CHEMISTRY.  [§  86. 

bers.  These  are  three  or  four  in  number  and  have  a  total  capacity 
of  4000-5000  cubic  meters.  Their  form  is  that  of  a  parallelepiped, 
whose  cross-section  is  nearly  a  square.  Lead  has  been  chosen  as 
the  material  for  the  walls  of  the  chambers,  because  it  is  the  only 
one  of  the  common  metals  which  is  only  slightly  attacked  by 
sulphuric  acid  and  the  substances  used  in  its  manufacture.  The 
lead  chambers  are  connected  with  each  other,  with  the  Glover 
tower  and  with  the  Gay-Lussac  tower  by  means  of  lead  pipes. 
The  first  two  chambers  are  also  furnished  with  openings  for 
introducing  steam. 

The  oxidation  of  dioxide  to  trioxide  having  been  accomplished 
in  the  lead  chambers,  the  residual  gas,  principally  nitrogen,  passes 
to  the  Gay-Lussac  tower. 

Usually  this  is  entirely  filled  with  coke.  On  top  of  the  tower 
is  a  reservoir  containing  60°-62°  sulphuric  acid  (BAUME,  see  §  88), 
which  comes  from  the  Glover  vtower.  The  Gay-Lussac  tower 
serves  to  collect  the  nitrous  vapors  that  are  still  present  in  the 
gas  as  it  leaves  the  lead  chambers.  These  vapors  dissolve  in  the 
sulphuric  acid,  forming  the  nitroso  sulphuric  acid  which  is  used 
in  the  Glover  tower.  In  this  way  the  loss  of  nitric  acid  is  much 
reduced. 

Let  us  now  examine  the  task  that  befalls  each  of  these  three 
— the  Glover  tower,  the  lead  chambers  and  the  Gay-Lussac 
tower. 

The  gases  that  come  from  the  pyrite  furnace  consist  of  a  mix- 
ture of  sulphur  dioxide  and  air,  a  larger  proportion  of  the  latter 
than  is  required  for  the  oxidation.  They  have  a  temperature  of 
about  300°  when  they  enter  the  Glover  tower,  A,  at  the  opening,  w. 
The  gas  current  rising  in  the  tower  meets  an  acid  mixture  flow- 
ing down  from  above.  The  latter  consists  of  the  nitroso  acid 
from  the  Gay-Lussac  tower,  diluted  with  the  acid  (chamber  acid) 

(3)  The  acid  produced  in  the  chambers  contains  about  67% 
H2S04  (53°  BAUME).  In  this  condition  it  is  employed  directly 
in  the  manufacture  of  fertilizers  ("superphosphate")-  For 
almost  all  other  purposes  it  must  ^rst  be  concentrated.  Ordinary 
sulphuric  acid  of  commerce  is  of  about  66°  B.  (B.  =  BAUME),  i.e. 
96-98%  H2S04.  It  is  prepared  from  the  chamber  acid  by  evap- 
orating it  first  in  lead  pans  to  about  78%  (60°  B.)  and  finally  in 
a  platinum  vessel, 


86.]  SULPHURIC   ACID.  -    137 

This  crude  sulphuric  acid  of  commerce  ("oil  of  vitriol")  still 
contains  various  impurities  and  is  usually  more  or  less  brown  in 
color  because  of  bits  of  straw  (from  the  packing  of  the  carboys) 
falling  in  and  charring.  It  can  be  purified  by  diluting  it,  where- 
upon the  dissolved  lead  sulphate  is  precipitated,  and  then  stirring 
in  a  little  barium  sulphide.  The  latter  produces  insoluble  barium 
sulphate,  and  also  hydrogen  sulphide,  which  precipitates  any 
arsenic  or  lead  (§  206)  still  present.  The  acid  is  then  decanted 
from  the  deposit,  concentrated,  and  finally  distilled. 

The  contact  process  . — It  has  already  been  stated  that 
sulphur  dioxide  unites  with  oxygen  directly  to  form  the  trioxide 
and  that  the  combination  is  considerably  accelerated  by  the  cata- 
lytic influence  of  platinized  asbestos.  This  simple  reaction  is  the 
basis  of  the  "  contact  process."  In  practice,  however,  air  is  used 
instead  of  pure  oxygen. 

The  process  falls  into  four  separate  parts:  1.  The  preparation 
of  a  mixture  of  sulphur  dioxide  and  air;  2.  The  purification  of  this 
mixture;  3.  The  formation  of  the  trioxide;  4.  The  combination 
of  sulphur  trioxide  with  water  to  form  sulphuric  acid. 

(1)  The  purification  of  the  gas  mixture  is  much  the  same  as  in 
the  lead-chamber  process.     For  reasons  which  will  soon  be  made 
clear  it  is  found  necessary  to  conduct  the  roasting  in  the  presence 
of  a  large  excess  of  oxygen.    While  the  equation 

2S02  +  O2=2SO3 

demands  only  1  vol.  O2  for  each  2  vols.  S02,  the  gases  are  usually- 
mixed  in  the  ratio  of  3  vols.  02  to  2  vols.  SO2. 

(2)  The    platinized    asbestos    acts   efficiently   only   when   the 
furnace  gases  are  absolutely  pure,  i.e.,  when  the  mixture  consists 
simply  of  sulphur  dioxide  and  air.     The  complete  purification  of 
these  gases  has  been  a  problem  of  exceptional  difficulty,  but  has 
been  accomplished  through  the  perseverance  of  KNIETSCH  of  the 
"Badische  Anilin-  und  Sodafabrik,"  the  great  chemical  factory  at 
Mannheim,  Germany.     In  the  first  place  the  furnace  gases  must  be 
wholly  freed  from  dust,  else  the  catalyzer  would  soon  become  so 
coated  as  to  lose  its  activity.     In  order  to  determine  when  the  gas  is 
really  dust-free  it  is  subjected  to  the  "optical  test,"  i.e.,  it  is 
passed  through  a  tube  closed  at  both  ends  with  glass,  and  is 


138  INORGANIC  CHEMISTRY.  [§S6- 

examined  with  the  eye  to  see  whether  it  is  perfectly  transparent 
and  free  from  nebulous  masses.  Even  when  this  optical  test  is 
quite  satisfactory  the  catalyzer  suffers  a  loss  in  activity  if  the 
gas  is  not  entirely  free  from  arsenic  compounds;  the  least  traces 
of  the  latter  have  an  injurious  effect.  The  presence  of  arsenic 
compounds  in  the  furnace  gas  is  due  to  the  occurrence  of  arsenic 
in  the  pyrites  (§  86;  1)  used  for  roasting.  KNIETSCH  has  finally 
succeeded  in  completely  eliminating  the  arsenic  compounds  by 
blowing  steam  into  the  gas  mixture. 

(3)   As  already  set  forth  in  §  79,  the  equilibrium 


+  O2<=±2SO3 
is  expressed  by  the  equation 


According  to  this  equation  the  formation  of  sulphur  trioxide  is 
more  complete  in  the  presence  of  an  excess  of  either  sulphur 
dioxide  or  oxygen,  for  as  p\  or  p2  increases  p3  must  also  increase. 
Since  the  object  in  view  is  to  convert  the  dioxide  as  completely  as 
possible  into  the  trioxide,  it  is  advantageous  to  provide  a  large 
excess  of  oxygen.  This  explains  why  more  than  the  theoretical 
amount  of  oxygen  is  taken.  Compare  (1). 

The  equilibrium  must  also  depend  on  the  pressure,  for,  if  this 
is  increased  n  times,  the  equation  becomes: 

or    np12p2 

from  which  it  is  evident  that  at  a  higher  pressure  (n>l)  the  for- 
mation of  the  trioxide  is  more  nearly  complete  (§  102,  5).  The 
manufacturer  does  not  find  it  necessary,  however,  to  employ 
high  pressure,  which  would  involve,  moreover,  a  great  complica- 
tion of  the  apparatus. 

If  it  is  desired  that  the  combination  of  sulphur  dioxide  and  oxy- 
gen should  be  as  complete  as  possible,  the  temperature  must  be 
kept  at  about  400°.  Since,  however,  the  heat  of  formation  of  the 
trioxide  is  great,  viz., 

SO2+0=S03+  22,600  Cal., 


86.]  SULPHURIC   ACID.  139 

the  apparatus  must  be  cooled.  This  is  done  most  practicably  by 
the  aid  of  a  fresh  portion  of  the  gas  mixture,  as  the  next  paragraph 
sets  forth. 

The  construction  of  the  apparatus  is  as  follows:  The  tubes  ab 
(Fig.  31)  contain  the  platinized  asbestos  6,  supported  on  little 
sieves  (shown  in  the  middle  tubes).  The  purified  furnace  gases 
first  pass  around  the  outside  of  the  tubes  and  are  thus  warmed  to 


FIG.  31. — CONTACT-PROCESS  APPARATUS. 

the  desired  temperature  at  the  heat  expense  of  the  gas  system 
within.  When  the  proper  temperature  is  reached  the  gases  are 
allowed  to  enter  the  tubes,  where  sulphur  trioxide  is  formed  with 
the  evolution  of  more  heat.  By  increasing  or  diminishing  the 
rate  of  flow  of  the  gas  current  the  temperature  can  be  regulated 
very  satisfactorily.  When  the  operation  is  started  the  apparatus 
must  first  be  warmed  to  400°. 

(4)  The  reaction  between  sulphur  trioxide  and  water  is  an 
energetic  one.     Nevertheless,  the  manufacture  of  sulphuric  acid 


140  INORGANIC    CHEMISTRY.  [§§86- 

from  these  two  compounds  involved  some  difficulty,  inasmuch  as 
sulphur  trioxide  fumes  invariably  escaped  when  this  substance  was 
introduced  into  water  or  dilute  sulphuric  acid.  Only  when  sul- 
phuric acid  of  97-98%  is  used  as  the  absorbent  and  care  is  taken 
to  keep  the  acid  at  this  concentration  by  the  simultaneous  addition 
of  water  does  a  complete  and  immediate  absorption  occur. 

This  is  due  to  two  circumstances:  first,  that  traces  of  water 
change  sulphur  trioxide  into  the  asbestine  modification  (§  80), 
which  is  only  slowly  absorbed  by  sulphuric  acid ;  second,  that  at 
the  concentration  of  97-98%  H2SO4  the  system  xS03+yR2O 
has  a  minimum  of  vapor  tension,  which  is  very  low. 

87.  Physical  Properties. — The  pure  compound,  hydrogen  sul- 
phate, is  an  oily  liquid  at  ordinary  temperatures,  solidifying  at  a 
low  temperature  and  melting  again  at  + 10.0°.  Its  specific  gravity 
in  the  liquid  state  (15°)  is  1.8500. 

Chemical  Properties. — The  concentrated  acid  obtained  by  dis- 
tillation is  not  the  simple  compound  H^SCU,  for  it  still  contains 
about  1.5%  of  water.  In  order  to  prepare  the  absolutely  pure  acid 
the  distilled  product  must  be  mixed  with  the  theoretical  amount 
of  sulphur  trioxide.  When  pure  sulphuric  acid  is  heated,  it  begins 
at  30°  to  give  off  fumes  of  sulphur  trioxide;  this  continues  until 
the  boiling-point,  317°  at  750  mm.  Hg.  pressure,  is  reached, 
when  an  acid  with  1.5%  water  distils  over.  On  heating  the  vapor 
of  sulphuric  acid  above  the  boiling-point,  it  begins  to  break  up 
into  water  and  the  anhydride;  this  dissociation  is  complete  at 
450°,  for  the  vapor  density  at  that  temperature  is  found  to  be  25.1, 
while  that  of  S03  +  H2O  is  theoretically  24.5. 

When  sulphuric  acid  is  mixed  with  water,  a  strong  evolution 
of  heat  occurs.  The  mixing  must  therefore  be  done  with  great 
care,  particularly  in  glass  vessels,  the  acid  being  poured  in  a  fine 
stream  into  the  water  and  the.  liquid  being  steadily  stirred.  On 
mixing  them  in  the  reverse  way,  by  pouring  the  water  into  the 
sulphuric  acid,  the  intense  heat  that  is  produced  may  cause  the 
glass  to  crack.  However,  when  the  acid  is  mixed  with  ice  in  a 
certain  proportion,  a  strong  cooling  follows. 

The  mixing  of  sulphuric  acid  and  water  is  attended  by  a  con- 
traction, i.e.  the  volume  of  the  dilute  acid  is  smaller  than  the  sum 
of  the  volumes  of  water  and  acid.  It  is  known  that  sulphuric 
acid  is  able  to  form  hydrates  with  water  (§  237). 


87.]  SULPHURIC  ACID.  141 

Sulphuric  acid  is  a  strong  dibasic  acid,  but  not  as  strong  as 
hydrochloric  acid,  for,  while  the  latter  is  ionized  to  95%  at  a 
dilution  of  0.1  gr.  mol.  per  L,  sulphuric  acid  at  the  same  dilution 
is  only  ionized  to  55%  into  2H'  +  S04".  At  higher  concentra- 
tions HSO4'  ions  also  exist.  It  acts  on  many  metals,  giving 
off  hydrogen.  This  action  is  made  use  of,  as  stated  above, 
in  the  preparation  of  hydrogen;  the  acid  must,  however,  be 
dilute,  for  when  it  is  too  strong  or  warmed,  the  hydrogen  gen- 
erated partially  reduces  the  sulphuric  acid  so  that  the  gas  given 
off  contains  hydrogen  sulphide.  Sulphur  dioxide  also  is  formed 
when  hydrogen  is  led  into  hot  sulphuric  acid.  It  is  upon  this 
action  that  the  reaction  of  copper  with  hot  concentrated  sulphuric 
acid  depends  (§  78).  Mercury,  silver  and  certain  other  metals  are 
similar  to  copper  in  their  behavior.  Platinum  and  gold  are  not 
attacked  by  the  acid. 

Sulphuric  acid  makes  holes  in  paper,  linen,  dress  goods  and 
the  like,  when  dropped  on  them.  It  has  a  destructive,  charring 
effect  on  organic  substances  in  general.  This  is  due  in  many 
cases  to  the  great  tendency  of  the  acid  to  unite  with  water,  which 
makes  it  not  only  deprive  other  substances  of  the  water  they  con- 
tain, but  even  withdraw  the  hydrogen  and  oxygen  from  organic 
compounds  to  form  water.  On  the  other  hand,  sulphuric  acid 
gives  up  oxygen  to  many  organic  substances,  being  itself 
reduced. 

In  order  to  detect  free  sulphuric  acid  in  vinegar,  for  example,  the 
liquid  is  evaporated  on  a  water-bath  with  a  little  sugar.  Free  sulphuric 
acid,  if  present,  chars  the  sugar  during  the  concentration. 

The  most  of  the  salts  of  sulphuric  acid  (sulphates) 
are  soluble  in  water.  Barium,  strontium,  and  lead  sulphates  are 
insoluble,  while  calcium  sulphate  (gypsum)  is  slightly  soluble,  but 
only  to  a  very  small  degree.  The  formation  of  barium  sulphate, 
BaS04,  serves  as  a  characteristic  test  for  sulphuric  acid,  or,  as  we 
may  better  say,  for  the  ion  $04". 

The  sulphates  are  in  general  very  stable.  They  can,  for 
instance,  be  heated  to  very  high  temperatures  without  decomposi- 
tion. The  acid  salts  lose  water  on  heating,  and  pass  over  into* 
pyrosulphates : 

2NaHSO4= H2O  +  Na2S2O7. 

Sodium 
nvrosulohate. 


142  INORGANIC    CHEMISTRY.  [§§  87- 

If  these  pyrosulphates  are  heated  still  higher,  they  give  off  sulphur 
trioxide  and  form  neutral  salts: 


88.  Uses.  —  Sulphuric  acid  is  of  enormous  practical  value,  its 
uses  being  most  varied.  It  is  employed  in  the  preparation  of  almost 
all  other  mineral  acids  from  their  salts.  In  the  manufacture  of  soda 
after  LE  BLANC  it  is  used  in  astonishingly  large  amounts  and  in 
nearly  all  other  branches  of  chemical  industry  it  is  of  some  service 
or  other.  In  the  laboratory  it  is  often  employed  as  a  drying-agent. 
A  moist  substance  is  dried  very  thoroughly  when  placed  in  a  closed 
apparatus  near  a  dish  of  the  concentrated  acid.  For  this  purpose 
special  pieces  of  apparatus  are  constructed,  called  desiccators. 

The  determination  of  the  concentration  of  sulphuric  acid  is  an 
operation  that  is  frequently  necessary.  Ordinarily  the  specific 
gravity  is  made  use  of,  for  this  can  be  determined  rapidly  with  an 
areometer.  There  are  tables  so  prepared  that  the  proportion  of 
H2SO4  or  80s  in  a  dilute  acid  whose  specific  gravity  and  tempera- 
ture are  known  can  be  quickly  read.  BAUME,  a  chemist  of  the 
latter  part  of  the  eighteenth  century,  constructed  an  areometer 
with  an  arbitrary  scale,  the  zero  point  of  which  indicates  pure 
water  and  the  point  10  being  reached  in  a  10%  salt-solution.  All 
the  divisions  are  equal.  100%  H2S04  would  then  be  represented 
by  the  line  66.6.  In  the  arts  the  strength  of  sulphuric  acid  is 
still  given  as  so  many  "  degrees  BAUME." 

Fuming  sulphuric  acid  is  the  name  of  a  sulphuric  acid  that 
contains  sulphur  trioxide  in  solution.  It  is  obtained  by  dissolving 
the  oxide  in  concentrated  sulphuric  acid. 

Fuming  sulphuric  acid  is  a  thick  oily  liquid,  which  fumes 
vigorously  in  the  air,  throwing  off  the  trioxide.  Sp.  g.  =  1.85-1.90. 

CHLORIDES  OF  SULPHURIC  ACID. 

89.  When  phosphorus  pentachloride  acts  on  sulphuric  acid  a  com- 
pound SO3HC1,  chlorosulphonic  acid,  is  formed  : 

H2S04  +  PC15  =  S03HC1  +  POC13  +  HC1. 

The  same  compound  results  from  the  direct  union  of  sulphur  trioxide 
and  hydrochloric  acid.     It  is  a  colorless  liquid,  which  fumes  vigorously  on 


91.]  CHLORIDES    OF   SULPHURIC  ACID.  143 

exposure  to  the  air.  Sp.g.=  1.766  at  18°.  Boiling-point,  158°.  On 
the  addition  of  water  a  violent  reaction  occurs,  producing  hydrochloric 
acid  and  sulphuric  acid: 

SO3HC1 +H2O=  H2SO4 +HC1. 

90.  A  compound,  SO2C12,  sulphuryl  chloride,  is  obtained  by  the  direct 
union  of  sulphur  dioxide  and  chlorine,  most  easily  by  first  saturating 
camphor  with  sulphur  dioxide  (which  readily  dissolves  in  it)  and  then 
passing  chlorine  over  it.  The  camphor  remains  unchanged.  Sulphuryl 
chloride  is  a  colorless  liquid,  which  boils  at  69.1°,  has  a  penetrating  odor, 
fumes  strongly  in  the  air  and  has  a  specific  gravity  of  1.6674  at  20°. 
The  addition  of  a  little  water  converts  it  into  chlorosulphonic  acid  and 
hydrochloric  acid,  much  water  to  sulphuric  and  hydrochloric  acids: 

SO2C12+  H2O=S03HC1+  HC1. 
SO,C12+2H2O=  H2SO4  +2HC1. 

These  decompositions  of  sulphuryl  chloride  can  be  represented  in  the 
following  way : 

OH 


In  the  place  of  the  two  chlorine  atoms  we  have,  therefore,  two  OH 
(hydroxyl)  groups  entering.  For  this  reason  it  is  assumed,  in  close 
analogy  with  the  methods  of  organic  chemistry,  that  sulphuric  acid  con- 
tains two  hydroxyl  groups. 

Sulphuryl  fluoride  can  be  obtained  by  the  direct  union  of  sulphur 
dioxide  and  fluorine.  It  has  the  same  remarkable  stability  as  the  com- 
pound SF8  (§  75).  It  is  a  colorless  and  odorless  gas,  liquid  at  —52°  and 
solid  at  —120°.  It  can  be  heated  with  water  in  a  sealed  tube  to  150° 
without  undergoing  decomposition.  Alkalies  absorb  it,  though  very 
slowly.  Sodium  can  be  fused  in  it  without  being  attacked. 

Persulphuric  Acid,  H2S2O8. 

91.  The  potassium  salt,  K2S2O8,  or,  still  better,  the  ammonium  salt, 
(NH4)2S208,  of  this  acid  can  be  obtained  by  the  electrolysis  of  a  cold 
saturated  solution  of  the  corresponding  sulphate  in  sulphuric  acid  of 
1.3  sp.  g.  In  such  a  solution  we  may  assume  we  have  the  ions  K*  and 
HSO/;  the  latter  are  discharged  at  the  anode  and  can  then  unite  to 
form  H2S2O8,  which  forms  with  the  K '  ions  present  the  difficultly  soluble 
potassium  salt  K2S2O8.  This  separates  out  as  a  white  crystalline  mass. 


144  INORGANIC  CHEMISTRY.  [§§91- 

However,  the  combination  of  two  HS04  groups  only  takes  place  when 
their  concentration  at  the  anode  is  quite  high;  for  if  this  is  not  the 
case  there  is  more  opportunity  for  secondary  reactions,  such  as  a  union 
with  water  to  form  2H2SO<  and  20H,  the  latter  of  which  is  decomposed 
into  H2O  and  O.  Such  a  high  concentration  at  the  anode  is  reached 
by  using  a  very  small  electrode.  The  electric  current  therefore  has  a 
high  density  at  the  anode;  that  is,  a  large  quantity  of  electricity  must 
pass  through  a  small  surface.  The  effect  thereof  is  that  this  large 
quantity  discharges  a  great  many  HSO/  ions  into  a  small  space,  or  in 
other  words,  produces  enough  HS04  groups  to  make  the  concentration 
ver}r  high  there. 

As  low  as  100°  it  decomposes  in  the  following  way: 

2K2S208=2K2S207+02. 

K-pyrosulphate. 

The  barium  salt  of  persulphuric  acid  is  soluble  in  water,  as  are  also 
most  of  the  other  known  salts. 

The  action  of  100%  hydrogen  peroxide  on  sulphur  trioxide  or  on 
chlorsulphonic  acid  yields  CARD'S  acid; 

SO3+H2O2  =  H2SO5, 

>OH  X)H 

S02  =  S02  +HC1. 

\C1+H2O2         \O-OH 

It  crystallizes  very  prettily  and  melts  at  about  45°  with  slight  decom- 
position. 

CARD'S  acid  reacts  with  another  molecule  of  chlorsulphonic  acid 
according  to  the  equation 


X)H  X)H       HO 

\O-OH 


XW-1.A  J.J.W 

SO2'  +C1-SO2OH=SO2  I 

\).0-S02 


forming  persulphuric  acid,  which  can  be  obtained  in  this  way  pure  and 
crystallized,  with  a  melting-point  of  60°  (attended  by  slight  decom- 
position). A  solution  of  CARD'S  acid  in  sulphuric  acid  can  be  prepared 
in  a  simple  way  by  mixing  H2O2  with  an  excess  of  strong  sulphuric  acid. 
On  the  basis  of  this  method  of  formation  BAEYER  gave  the  compound  the 
name  sulpho-mono-peracid. 

It    has    very  strong  oxidizing  powers.      It  sets  iodine  free   from 
potassium  iodide,  oxidizes  sulphur  dioxide  to  trioxide,  and  ferrous  to 


92.]  POLYTHIONIC  ACIDS.  145 

ferric  salts  and  also  precipitates  the  higher  oxides  of  silver,  copper, 
manganese,  cobalt,  and  nickel  from  solutions  of  salts  of  these  metals. 
On  the  other  hand,  it  neither  bleaches  permanganate  solution  nor 
oxidizes  solutions  of  chromic  and  titanic  acids;  in  these  respects  it  is 
distinguished  from  hydrogen  peroxide,  to  which  it  otherwise  shows 
much  similarity. 

POLYTHIONIC  ACIDS. 

92.  Under  this  name  are  grouped  four  acids  of  the  general 
formula  H2SnO6,  in  which  the  number  of  sulphur  atoms,  n,  can 
be  2,  3,  4  and  5,  and  this  determines  the  names  of  the  individual 
acids. 

Dithionic  acid,  H2S206.  The  manganese  salt  of  this  acid  is  obtained 
when  finely  powdered  manganese  dioxide  is  suspended  in  water  and  sul- 
phurous oxide  passed  in: 

2S02+Mn02=MnS2O6. 

From  this  barium  salt  the  dithionic  acid  can  be  liberated  by  sulphuric 
acid.  The  solution  can  ba  concentrated  in  vacuo  over  sulphuric  acid  till 
its  specific  gravity  reaches  1.347;  farther  concentration  or  warming 
results  in  a  decomposition: 


Trithionic  acid,  H2S3O6.  Potassium  trithionate  is  formed  when  a 
solution  of  potassium  thiosulphate  is  saturated  with  sulphur  dioxide: 

3SO2  +2K2SA  =  2K2S3O6+  S. 

The  free  acid  is  unstable  ;  even  at  ordinary  temperatures  it  decomposes 
in  a  dilute  solution  into  sulphur,  sulphurous  oxide  and  sulphuric  acid  : 

H2S3O6=H2S04+S+S02. 

Tetrathionic  acid.  Its  salts  result  from  the  action  of  iodine  on  the  solu- 
tion of  a  thiosulphate. 

K2S2O3+  21  -»2KI+  K2S4O6. 

The  acid  itself  can  be  obtained  (also  only  in  dilute  solution)  by  adding 
sulphuric  acid  to  the  barium  salt,  which  is  prepared  in  an  analogous 
manner.  In  dilute  solution  it  is  quite  stable;  in  the  concentrated  state 
it  breaks  up  into  sulphur,  sulphurous  oxide  and  sulphuric  acid. 

Pentathionic  acid.  On  mixing  solutions  of  sulphur  dioxide  and 
hydrogen  sulphide  the  principal  reaction  is  a  mutual  oxidation  and 
reduction  of  these  compounds  with  the  separation  of  sulphur  (§  78). 
The  action  is,  however,  much  more  complicated,  inasmuch  as  polythionic 
acids,  among  them  pentathionic  acid,  are  formed  in  addition  at  the  same 
time.  The  mixture  of  H2S.aq  and  SO2.aq  is  known  as  "  WACKENRODER'S 
liquid."  Well-crystallized  salts  of  pentathionic  acid  have  been  prepared. 


146  INORGANIC  CHEMISTRY.  [§93. 

Use   of   Sodium    Thiosulphate     in   Volumetric 
Analysis.     lodometry. 

93.  On  adding  sodium  thiosulphate  to  an  iodine  solution,  the 
intensely  brown  liquid  loses  its  color,  sodium  iodide  and  sodium 
tetrathionate,  two  colorless  compounds,  being  formed: 

2Na2S2O3 + 21  =  Na2S4O6 + 2NaI  ; 
or,  writing  only  the  ions  that  take  part  in  the  reaction: 


The  disappearance  of  the  color  is  thus  due  to  the  fact  that  the 
molecules  of  iodine  are  transformed  into  ions  by  taking  up.  two 
negative  charges  from  2S2(V.  Upon  this  fact  a  method  is  based 
for  determining  the  amount  of  free  iodine  in  a  solution.  This  is  done 
by  allowing  a  solution  of  sodium  thiosulphate,  whose  concentration 
(litre)  is  known,  to  flow  drop  by  drop  into  a  definite  volume  of  iodine 
solution.  (For  letting  out  a  certain  amount  of  liquid  a  pipette 
(Fig.  32)  is  commonly  employed.)  The  color  gradually  brightens 
and  finally  a  point  is  reached  when  the  liquid  is  only  slightly  tinged 
and  the  addition  of  another  drop  causes  the  color  to  entirely  dis- 
appear. This  transition  can  be  very  accurately  detected.  The 
iodine  molecules  have  now  entirely  disappeared.  Since  according 
to  the  above  equation  a  molecule  of  thiosulphate  is  consumed  for 
each  atom  of  iodine,  the  percentage  of  iodine  in  the  solution  can 
be  calculated  from  the  amount  of  thiosulphate  used. 

To  make  the  calculation  of  the  result  of  such  a  determination 
(titration)  as  easy  as  possible  the  thiosulphate  solution  is  so  stand- 
ardized that  it  bears  a  certain  relation  to  an  equivalent  of  iodine 
(  =  127  g.),  i.e.  a  certain  amount  bleaches  exactly  this  much  iodine. 

"Normal  solution  "  is  a  name  applied  to  a  solution  containing 
the  equivalent  weight  (§  23)  in  grams  (gram  equivalent)  in  one 
liter.  Frequently  use  is  also  made  of  a  £,  J,  TV  or  a  twice,  thrice, 
etc.,  normal  solution.  Normal  hydrochloric  acid  contains  36.5  g. 
HC1,  normal  sulphuric  acid  49  g.  H2SC>4  (=  i  gram  molecule), 
a  normal  iodine  solution  127  g.  iodine,  per  liter.  Detailed  direc- 


§  93.] 


VOLUMETRIC   ANALYSIS. 


147 


tions  for  preparing  such  solutions  can  be  found  in  the  text-books 
of  analytical  chemistry. 

In  order  to  determine  readily  the  volume  of  thiosulphate  solu- 
tion that  is  required  in  the  analysis,  use  is  made  of  a  burette 
(Fig.  33),  a  glass  tube  that  is  divided  into  -^  c.c.  and  closed  at 
the  lower  end  with  a  glass  stop-cock  or  with  a  rubber  tube  and 
pinch-clamp.  In  titrating  the  iodine  solution  the  thiosulphate 


A 


FIG.  32.— PIPETTE. 


FIG.  33.  — BURETTES  AND  SUPPORT. 


solution  is  allowed  to  flow  out  slowly  and,  finally,  drop  by  drop, 
while  the  liquid  is  being  stirred. 

Example.  For  50  c.c.  of  an  iodine  solution  whose  strength  is 
to  be  determined  27.30  c.c.  T\  normal  thiosulphate  solution  was 
necessary  before  the  color  completely  disappeared.  Required  the 
number  of  grams  of  iodine  contained  in  1  liter  of  this  solution. 

1000  c.c.  TV  normal  Na2S2O3  solution  (see  above)  decolorizes 
TV  equivalent  of  iodine  (  =  12.7  g.);  27.3  c.c.  therefore  decolorizes 

12  7 
27.3Xinnn   g.  iodine.    This  amount  is  contained  in  50  c.c.  of 

lUUU 


148  INORGANIC    CHEMISTRY.  [§§93- 

the  iodine  solution  in  question.     Hence  1  liter  of  the  latter  con- 
tains 20  X 27.3  X  12.7  X  10-3=6.8842  g.  iodine. 

Various  other  substances  which  liberate  iodine  from  potassium 
iodide  can  be  determined  by  titrating  the  amount  of  iodine  dis- 
placed; for  example,  chlorine  and  bromine  may  be  thus  determinedf 
since  they  set  free  the  equivalent  amount  of  iodine  from  potassium 
iodide  solution. 

SELENIUM. 

94.  Selenium  was  discovered  by  BERZELIUS  in  1817.  It 
took  its  name  from  (rekrjvrj  (the  moon) ,  because  it  possesses  great 
similarity  to  the  element  tellurium  (named  from  tellus=ihe  earth) 
discovered  a  short  time  previously.  It  is  rather  widely  distributed 
in  nature,  but  it  occurs  only  in  small  quantities.  It  is  found 
native,  is  frequently  found  in  pyrite  and  also  appears  in  some 
rare  minerals.  When  this  sort  of  pyrite  is  employed  in  sulphuric 
acid  manufacture,  the  selenium  collects  in  the  "  chamber-mud  " 
of  the  lead  chambers;  from  this  it  is  usually  obtained. 

The  process  is  as  follows:  The  selenium  deposit  is  heated  with  nitric 
acid,  which  oxidizes  the  selenium  to  selenic  acid,  H2SeO4.  The  solution 
thus  obtained  is  first  boiled  with  hydrochloric  acid,  whereby  selenious 
acid,  H2Se03,  is  formed  with  the  evolution  of  chlorine.  This  latter  acid 
is  then  reduced  by  means  of  sulphurous  oxide  to  selenium,  which  separates 
in  amorphous  red  flakes. 

Selenium  displays  analogy  with  sulphur  in  many  respects;  for 
instance,  in  occurring. in  various  allotropic  conditions.  According 
to  SAUNDERS,  there  is  an  amorphous  red  modification,  that  is 
soluble  in  carbon  disulphide.  From  this  solution  the  selenium 
separates  as  a  second  modification,  which  is  the  red  crystalline 
selenium,  fusing  at  170°-180°.  Then  there  is  a  metallic  form  fusing 
at  217°.  This  modification  appears  when  amorphous  selenium  is 
heated  to  97°,  at  which  point  a  sudden  and  marked  rise  of  tempera- 
ture occurs;  or  when  molten  selenium  is  suddenly  cooled  to  210° 
and  kept  for  a  time  at  that  temperature.  In  this  metallic  state 
selenium  has  a  metallic  lustre,  is  insoluble  in  carbon  disulphide  and 
conducts  electricity.  Its  conductivity  strangely  depends  very 
much  on  the  intensity  of  its  illumination,  however. 

The  melting-point  of  selenium  is  217°,  its  boiling-point  680°. 
As  in  the  case  of  sulphur  the  vapor  density  decreases  with  rising 


94.]  SELENIUM.  149 

temperature  till  about  1400°  is  reached,  when  it  remains  constant. 
At  this  temperature  it  is  found  to  be  81.5  (H=l),  corresponding 
to  a  molecular  weight  of  163.0.  Now  since  the  atomic  weight  of 
selenium,  as  deduced  from  the  vapor  density  of  its  compounds, 
is  78.9,  the  above  molecular  weight  agrees  very  closely  with  the 
formula  862. 

Hydrogen  selenide,  H2Se,  can  be  obtained  directly  from  its 
elements,  as  these  unite  at  400°.  Analogously  to  hydrogen  sul- 
phide, it  can  also  be  got  by  the  decomposition  of  iron  selenide, 
FeSe,  with  hydrochloric  acid.  At  a  high  temperature  hydrogen 
selenide  dissociates  into  its  elements.  Its  properties  are  only 
slightly  acidic  and  it  is  more  poisonous  than  sulphuretted  hydrogen. 
The  heavy  metals  are  precipitated  from  their  solutions  as  selenides 
by  it. 

An  aqueous  hydrogen  selenide  solution  becomes  turbid  on 
standing  because  of  the  selenium  that  separates  out. 

Two  chlorine  compounds,  Se2Cl2  and  SeCU,  are  known.  The 
latter  is  much  more  stable  than  the  corresponding  sulphur  com- 
pound, SOU  (§  75).  Selenium  tetrachloride  is  solid  and  sublimes 
without  decomposition;  dissociation  does  not  begin  until  200°  is 
reached. 

Selenium  dioxide,  Se02,  is  the  only  oxide  of  selenium  known. 
It  results  from  the  burning  of  selenium  in  the  air.  The  extremely 
disagreeable  odor  which  arises  is  not  a  property  of  the  dioxide, 
however,  but  is  probably  due  to  the  formation  of  another  oxygen 
compound  of  selenium  which  has  not  as  yet  been  isolated.  Sele- 
nium dioxide  forms  long  white  needles  that  sublime  at  310°. 

Selenium  dioxide  is  an  acid  anhydride;  on  dissolving  it  in  water 
an  acid,  selenious  acid,  H2SeO3,  is  formed,  which  can  be  isolated 
(unlike  sulphurous  acid).  This  acid  crystallizes  in  large  colorless 
prisms.  On  being  heated  it  breaks  up  into  water  and  anhydride. 
Sulphur  dioxide  or  stannous  chloride  reduce  it  to  free  selenium, 
which  is  deposited  in  red  flakes: 

H2Se03 + 2SO2  +  H20  =  2H2SO4 + Se. 

Sulphuretted  hydrogen  precipitates  from  the  solution  selenium 
sulphide,  SeS,  insoluble  in  ammonium  sulphide. 

When  chlorine  is  passed  into  the  solution  of  selenious  acid  or 


150  INORGANIC  CHEMISTRY.  [§§94- 

when  bromine  is  added  to  it,  selenic  acid,  H2Se04,  is  formed.  In 
the  pure  state  this  is  a  crystalline  solid,  melting  at  58°.  The  95% 
solution  of  it  is  an  oily  liquid,  which  has  the  appearance  of  sul- 
phuric acid.  The  barium  salt  of  the  acid,  like  that  of  sulphuric 
acid,  is  extremely  difficultly  soluble. 

On  boiling  with  hydrochloric  acid,  selenic  acid  is  reduced  to 
selenious  acid  with  the  evolution  of  chlorine. 

Tellurium. 

95.  Tellurium  is  of  rare  occurrence;  it  is  known  in  the  rative  condi- 
tion and  also  in  combination  with  bismuth,  and  with  gold  or  silver  (in 
sylvanite,  or  graphic  tellurium}.  It  is  found  chiefly  in  Transylvania  and 
in  the  Altai  mountains,  and  also  in  Boulder  Co.,  Colorado.  In  the  amor- 
phous condition  tellurium  is  a  black  powder,  but  after  fusion  it  is  silvery 
white,  of  a  metallic  lustre  and  a  conductor  of  heat  and  electricity.  The 
vapor  density,  as  in  the  cases  of  selenium  and  sulphur,  decreases  with 
increasing  temperature  and  does  not  remain  constant  till  about  1400°; 
it  then  corresponds  to  a  Te2  molecule. 

Hydrogen  telluride,  H2Te,  results  from  the  action  of  hydrochloric  acid 
on  zinc  telluride,  ZnTe.  The  product  thus  obtained  contains  more  or 
less  hydrogen.  It  is  very  poisonous,  and  dissociates  readily.  From  solu- 
tions of  the  heavy  metals  it  precipitates  their  tellurium  compounds 
(telluride  s). 

Tellurium  dioxide,  TeO2,  is  formed  on  burning  tellurium  in  the  air. 
It  is  very  difficultly  soluble  in  water. 

Tellurous  acid,  H2TeO3,  is  obtained  by  dissolving  tellurium  in  nitric 
acid.  It  dissolves  in  water  with  great  difficulty  and  breaks  up  on  warming 
into  TeO2  and  H2O. 

Telluric  acid,  H2Te04,  is  prepared  by  fusing  the  metal  or  the  dioxide 
with  soda  and  saltpetre  and  separating  the  acid  from  the  tellurate  formed. 
The  compound  H2Te04  +  2?I2O  crystallizes  out  from  the  aqueous  solution; 
it  loses  its  water  of  crystallization  at  100°.  The  free  telluric  acid,  H2TeO4, 
prepared  in  this  way  is  a  white  powder,  difficultly  soluble  in  cold  water. 
Telluric  acid  has  only  feebly  acid  properties. 

Selenium  and  tellurium  both  combine  with  potassium  cyanide,  when 
they  are  fused  with  it,  forming  compounds  corresponding  to  KCNS,  viz., 
KCNSe  and  KCNTe.  Nevertheless,  while  potassium  t  e  1 1  u  r  o-cyanide 
is  at  once  decomposed  by  the  oxygen  of  the  air  with  the  separation  of 
tellurium,  potassium  s  e  1  e  n  i  o-cyanide  is  more  stable  and  does  not 
decompose  with  the  separation  of  selenium  until  it  is  boiled  with  hydro- 


96.] 


SUMMARY  OF  THE  OXYGEN    GROUP. 


151 


chloric  acid.    We  have  here  a  means  of  detecting  selenium  in  the  pres- 
ence of  tellurium  and  of  separating  the  two. 


SUMMARY  OF  THE  OXYGEN  GROUP. 

96.  The  elements  oxygen,  sulphur,  selenium  and  tellurium, 
like  the  halogens,  form  a  natural  group,  particularly  in  two  respects; 
their  compounds  correspond  to  a  general  type  and  their  physical 
and  chemical  properties  vary  gradually  with  increasing  atomic 
weight.  Their  hydrogen  compounds  have  the  formula  RH2,  their 
oxygen  compounds  and  their  acids  the  formulae  RO2  and  H2RO3, 
and  also  ROs  and  H2RC>4.  Ozone  may  be  considered  with  reference 
to  these  types  as  analogous  to  sulphur  dioxide;  O-O2  ozone;  8-62 
sulphur  dioxide. 

The  following  table  shows  the  gradual  change,  or  progression, 
of  the  physical  properties : 


O. 

S. 

Se. 

Te. 

Atomic  weight  
Specific  gravity.  .  .  . 

Melting-point  . 

16.00 
1.124 

(at  -181°) 

32.07 
1.95-2.07 

119.5° 

79.2 
4.2-4.8 

217° 

127.5 
6.2 

452° 

Boiling-point  . 

—  181  4° 

450° 

680° 

white  heat 

Color 

light  blue 

yellow 

red 

black 

As  the  atomic  weight  increases,  the  values  of  the  physical  con- 
stants also  increase,  as  the  table  shows.  At  the  same  time  the 
external  appearance  approaches  that  of  the  metals;  in  tellurium 
the  metallic  appearance  is  quite  marked. 

The  instability  of  the  hydrogen  compounds  increases  from  oxygen 
to  tellurium;  the  strength  of  the  oxygen  acids  diminishes  rapidly, 
sulphuric  acid  belonging  to  the  strongest,  and  telluric  acid  to  the 
very  weak,  acids. 

It  should  also  be  noted  that  all  of  these  elements  appear  in 
allotropic  modifications. 


152  INORGANIC    CHEMISTRY.  [§§97- 

THERMOCHEMISTRY. 

97.  It  was  stated  above  (§  20)  that  a  chemical  combination  or 
decomposition  is  accompanied  by  an  evolution  or  absorption  of 
heat,  in  other  words  by  a  heat  change,  or  caloric  effect.  In  many 
cases  this  caloric  effect  has  been  carefully  measured.  The  work 
of  BERTHELOT  and  of  THOMSEN  along  this  line  has  been  especially 
fruitful.  That  part  of  chemistry  which  deals  particularly  with 
these  caloric  effects  is  called  thermochemistry. 

The  caloric  effect  is  always  given  for  molecular  amounts  of  the 
reacting  substances,  since  in  this  way  only  is  it  possible  to  compare 
substances  from  a  chemical  standpoint.  Hence,  when  the  heat  of 
formation  of  water  is  said  to  be  69.0  calories  (kilogram  calories), 
it  is  implied  that  this  number  of  calories  is  evolved  by  the  union 
of  2  g.  hydrogen  with  16  g.  oxygen: 

2H  +  O  =  H2O  +  69.0  Cal. 

In  this  equation  H  and  O  stand  for  gram  atoms. 

In  expressing  a  caloric  effect  it  is  necessary  to  indicate  the  state 
of  matter  of  the  reacting  and  the  resulting  substances,  in  so 
far  as  this  is  not  self-evident,  because  the  latent  heat  of  fusion 
or  vaporization  must  be  taken  into  consideration.  The  above 

amount, 

69.0  Cal., 


refers  to  the  formation  of  water  and  its  conversion  to  a  liquid.  It 
therefore  includes  the  heat  of  condensation.  Since  this  amounts 
to  0.536  Cal.  per  gram,  it  would  in  this  case  (for  18  g.)  be  9.6  Cal.; 
hence  the  caloric  effect  of  the  combustion  of  hydrogen  to  steam  at 

100°  is 

2H  +  0  =  H2Ogas  +  58.4  Cal. 

The  caloric  effect  is  also  influenced  by  the  state  of  matter 
in  which  the  substances  react,  i.e.,  whether  solid,  liquid,  or  gas, 
inasmuch  as  solution  is  almost  always  accompanied  by  a  heat 
change.  In  the  formation  of  sodium  chloride  by  the  mixture  of 
dilute  solutions  of  sodium  hydroxide  and  hydrochloric  acid  (this 
being  indicated  by  aq  after  the  formulae  of  the  substances)  the 
caloric  effect  is: 

NaOHaq  +  HClaq  =  NaClaq  +  H2O  +  13.7  Cal. 


99.]  THERMOCHEMISTRY.  153 

However,  when  the  salt  is  prepared  by  passing  hydrochloric 
acid  gas  into  a  dilute  solution  of  the  base,  the  equation  is  as  follows: 

NaOHaq+HClgas  =NaClaq  +  H2O  +  31.1  Cal. 

We  thus  obtain  13.7  Cal.  as  before,  but  increased  by  the  heat  of 
solution  of  gaseous  hydrochloric  acid  in  a  large  amount  of  water, 
viz.,  17.4  Cal. 

The  heat  of  formation  of  chemical  compounds  must  be  equal 
to  their  heat  of  decomposition,  but  have  the  opposite  sign.  Were 
this  not  the  case,  heat  would  be  lost  or  gained  when  a  compound 
is  formed  and  then  decomposed  so  as  to  return  to  the  original  con- 
dition, and  such  a  result  would  be  at  variance  with  the  Law  of  the 
Conservation  of  Energy. 

Experience  has  shown  that  in  the  formation  of  most  compounds 
heat  is  generated,  but  that  in  many  cases  heat  is  absorbed.  Chem- 
ical actions  of  the  first  sort  are  called  exothermic,  those  of  the 
second  endothermic,  reactions.  An  example  of  the  second  sort 
is  the  synthesis  of  chlorine  monoxide: 

2Cl  +  0  =  Cl2Ogas-15.1  Cal. 

98.  For  the  determination  of  the  caloric  effect  various  methods  are  in 
use.     Only  those  actions  are  suitable  for  thermochemical  measurements 
which  complete  themselves  quickly.     In  measuring  the  caloric  effect  in 
the  case  of  liquids  or  solutions,  as,  for  example,  the  heat  of  neutralization 
of  acids  and  bases,  the  heat  of  solution  or  of  dilution,  etc.,  an  ordinary 
calorimeter  is  generally  used,  such  as  is  employed  in  physics  for  the 
method  of  mixtures,  the  same  precautions  being  taken  in  order  to  secure 
accurate  results. 

The  heat  of  combustion  of  a  substance  is  usually  measured  with  the 
calorimetric  bomb  of  BERTHELOT-MAHLER.  This  is  the  usual  method  with 
organic  compounds. 

99.  The  Law  of  HESS.     The  entire  caloric  effect   (the   whole 
amount  of  energy)  produced  by  the  transformation  of  one  chemical 
system  into  another  is  independent  of  all  intermediate  stages. 

This  law  is  a  direct  consequence  of  the  principle  of  the  con- 
servation of  energy.  If  HESS'S  law  did  not  hold,  energy  would 
have  to  be  gained  or  lost  in  the  transition  from  one  system  to 
another  and  the  subsequent  return  to  the  initial  condition,  which 


154  INORGANIC  CHEMISTRY. 


[§99- 


is   contradictory  to  the  above  principle.     A  few  examples  will 
serve  to  make  this  law  more  clearly  understood. 

(a)  A  dilute  solution  of  sodium  sulphate  can  be  prepared 
from  sodium  hydroxide,  sulphuric  acid  and  water  in  various  ways. 
For  instance,  two  gram-molecules  of  the  base  can  be  treated  at 
once  with  dilute  sulphuric  acid;  or  one  gram-molecule  of  the 
base  can  be  mixed  with  the  acid  at  first  and  the  second  added 
afterward.  Accordingly  we  get  the  following  caloric  effects: 

(1)     2NaOH  +H2S04aq     -Na2SO4aq  -2H2O     =  31.4Cal. 

,     C  NaOHaq  +  H2SO4aq       -  NaHSO4aq  -  H2O          =        14 .  75 

(  }  (  NaOHaq + NaHSO4aq  -  Na2S04aq  -  H20      =     1 6 . 65 


Total...  ..........................  31.4Cal. 

(6)  From  ammonia,  hydrogen  chloride  and  water  a  dilute  solu- 
tion of  ammonium  chloride,  NH4C1,  can  be  prepared,  either  by 
letting  dry  ammonia  gas  combine  with  dry  hydrogen  chloride  gas 
and  dissolving  the  resulting  ammonium  chloride  in  water  or  by 
first  dissolving  ammonia  and  hydrogen  chloride  in  separate  por- 
tions of  water  and  then  mixing  the  solutions.  In  the  first  case  we 
have  the  equations: 

NH3gas  +  HClgas  -  NH4Clsoiid  ............    =     42  .  6 

NH4Clsoiid  +  aq-NH4Claq  ..............    =  -  4.0 

38.6Cal. 
in  the  second  case: 


-NH3aq  ...............    =  8.82 

HCl  +  aq-HClaq  ................    =17.13 

NH3aq  +  HClaq  -  NH4Claq  ........     =  12  .  45 

38.40Cal. 

The  final  effects  in  the  two  cases  are  found  to  be  alike  within  the 
limits  of  experimental  error. 

With  the  help  of  HESS'S  law  the  determination  of  the  caloric 
effect  is  rendered  possible  in  many  reactions  which  cannot  be 
dealt  with  directly  or  are  unsuitable  for  calorimetric  measure- 
ments. In  general  this  is  done  by  making  thermochemical  meas- 


100.]  THERMOCHEMISTRY.  155 

urements  for  a  series  of  processes  in  which  the  reaction  plays  a 
part  and  finally  calculating  the  caloric  effect  of  the  reaction  as 
the  single  unknown,  as  will  be  more  fully  explained  in  the  examples 
below. 

Suppose  it  were  required  to  find  the  heat  of  formation  of  hydro- 
gen sulphide.  This  compound  can  be  formed  directly  from  its 
elements  (§  72),  but  the  reaction  is  unsuitable  for  thermochemical 
study.  We  will  therefore  start  with  the  system,  H,  S,  and  O,  and 
consider  the  two  ways  by  which  it  can  form  water  and  sulphur 
dioxide:  (1)  hydrogen  and  sulphur  are  burned  directly  to  water 
and  sulphur  dioxide;  (2)  (a)  hydrogen  and  sulphur  combine  and 
(6)  the  resulting  hydrogen  sulphide  is  burned  to  water  and  sulphur 
-dioxide.  Since  we  started  with  the  same  system  and  in  the  end 
reached  the  same  result  in  each  case,  the  caloric  effect  must  be  the 
same  according  to  HESS'S  law,  so  that,  if  we  measure  (1)  and  (26), 
we  can  equate  (1)  and  (2)  and  solve  for  (2a),  thus: 

Heat  of  combustion  of  2H-J-heat  of  combustion  of  S  = 
heat  of  formation  of  H2S  +  heat  of  combustion  of  H2S. 
(2H  +  0  -H20)  +  (S  +  20  -S02)  = 
(2H+S  -H2S)  +  (H2S  + 30  -S02  -H20). 
68.0  +  69.26  =  x+ 133.46; 
.;  %    z=(S+2H-H2S)  =  3.8. 

100.  In  using  these  values  of  the  heat  of  formation  and  heat  of 
decomposition  it  should  be  noted  that  they  do  not  represent  the 
amounts  of  heat  liberated  by  the  combination  of  atoms  to  form 
molecules,  but  that  the  heat  of  decomposition  of  the  molecules 
of  the  elements  (i.e.  the  amount  of  heat  required  to  break  these 
molecules  up  into  atoms)  is  always  included.  When,  for  example, 
chlorine  unites  with  hydrogen  to  form  hydrochloric  acid,  22.0  Cal. 
are  given  off.  That  which  is  measured  is  the  total  caloric  differ- 
ence between  the  initial  system  H2+Cl2  and  the  2HC1  formed 
from  it.  In  the  indirect  determination  of  a  heat  of  formation 
with  the  help  of  HESS'S  law  the  calculated  caloric  effect  also 
includes  the  heat  of  decomposition  of  the  molecules  of  the  ele- 
ments. In  the  determination  of  the  heat  of  formation  of  hvdrogen 


156  INORGANIC    CHEMISTRY.  [§§100- 

sulphide,  for  instance,  in  the  above  way  the  caloric  effect  of  the 
combustion  of  this  gas  is  composed  of  the  following  parts: 

2(2H  +  S  -  H2S)  +  3(2O  -  O2)  =  2SO2 + 2H2O  +  p  Cal. ; 
that  of  the  combustion  of  hydrogen  of  the  following: 
2(2H  -  H2)  +  (20  -  O2)  =  2H2O  +  q  Cal. ; 
that  of  the  combustion  of  sulphur  of 

(2S-S2)+2(2O-O2)  =  2SO2+rCal.; 

(20  — 02),  etc.,  indicating  the  heat  of  decomposition  of  molecules 
of  the  elements. 

The  heat  of  formation  of  hydrogen  sulphide  is  r+q—p.  Deduc- 
ing the  value  of  r+q—p  from  the  above  equations,  we  have 

r+g-p=(2S-S2)+2(2H-H2)-2(2H-fS-H2S), 

from  which  it  follows  that  the  heats  of  formation  of  the  sulphur 
and  hydrogen  molecules  are  included  in  the  heat  of  formation 
found. 

CHEMICAL  AFFINITY. 

101.  When  a  compound  is  formed,  we  attribute  the  phe- 
nomenon to  the  affinity  which  exists  between  the  combining  sub- 
stances. The  term  "affinity"  comes  down  from  an  age  when 
it  was  thought  that  only  those  substances  could  combine  with 
one  another  which  were  in  a  certain  agreement  with  each  other 
(were  "  in  love  with  each  other,"  as  EMPEDOCLES  and  later  also 
GLAUBER  expressed  it) . 

This  affinity  was  originally  considered  as  a  force.  THOMSEN, 
for  example,  defined  it  as  the  force  which  holds  the  parts  of  a 
compound  together.  Concerning  the  magnitude  of  this  force 
our  knowledge  was  for  a  long  time  only  qualitative.  If  the 
substances  AB  and  C  interacted  to  form  AC  and  B}  it  was  said 
that  the  affinity  of  A  for  C  was  greater  than  that  of  A  for  B. 
Comparative  study  of  such  reactions  led  to  the  arrangement  of 
a  series  of  the  elements  in  decreasing  order  of  affinity;  but  the 
absolute,  or  even  relative,  magnitude  of  these  affinities  was 
as  it  were  a  closed  book.  Henco  it  was  a  great  step  forward  when 


§101.]  THERMOCHEMISTRY.  157 

BERTHELOT  developed  a  method  of  measuring  affinity.  He  con- 
sidered that  the  quantity  of  heat  liberated  in  the  formation  of  a 
chemical  compound  was  a  measure  of  the  affinity  satisfied  by  the 
action.  Thus  affinity  came  to  be  regarded  no  longer  as  a  force, 
but  as  an  amount  of  work.  We  know  that  when  water  is  decom- 
posed by  the  current  from  a  dynamo,  work  must  be  done  in  order 
to  drive  the  dynamo  and  also  to  split  up  the  water  molecules; 
and,  conversely,  when  hydrogen  and  oxygen  unite,  heat,  or  in 
other  words  energy,  is  produced.  A  mixture  of  hydrogen  and 
oxygen  can  be  compared  with  a  lifted  stone ;  both  possess  potential 
energy.  When  the  stone  falls,  its  potential  energy  is  transformed 
into  kinetic  energy.  When  hydrogen  combines  with  oxygen  the 
potential  energy  of  the  system  is  converted  into  heat.  Since  he 
regarded  this  heat  effect  as  a  measure  of  the  driving  force  of  any 
chemical  reaction,  BERTHELOT  was  led  to  propose  his  principe 
du  travail  maximum,  viz.,  that  of  all  the  chemical  processes  which 
can  proceed  without  the  application  of  energy  from  an  outside 
source  that  one  always  occurs  which  involves  the  greatest  evolu- 
tion of  heat. 

However,  this  principle  did  not  prove  to  be  universally  applic- 
able. The  very  existence  of  endothermic  compounds  is  at 
variance  with  it,  for  the  heat  effect  of  a  reaction  involving  an 
endothermic  compound  would  be  greater  if  that  compound  were 
not  formed.  Further,  the  rapidly  increasing  number  of  known 
equilibrium  reactions  throws  doubt  on  the  principle,  for,  if  in  an 
equilibrium  A+B<=±AB  the  direct  reaction  (— »)  is  exothermic, 
the  opposing  reaction  (<— )  must  be  endothermic. 

Yet,  even  though  the  principle  could  not  be  accepted  as  a 
general  truth,  chemists  had  to  admit  that  in  very  many  cases 
it  represented  the  facts,  that  is,  it  contained  a  considerable 
amount  of  truth. 

VAN'T  HOFF  succeeded  in  putting  things  in  their  proper 
light.  The  amount  of  heat  liberated  in  a  chemical  reaction 
represents  the  total  change  of  the  energy  of  the  system,  and  this 
is  what  BERTHELOT  regarded  as  a  measure  of  the  affinity.  VAN'T 
HOFF  rejected  this  notion  and  showed  that  it  is  the  "free  energy" 
gained  in  a  reaction  which  must  be  regarded  as  a  measure  of  the 
affinity.  By  "free  energy"  we  understand  the  greatest  amount 
of  work  which  the  reaction  is  capable  of  doing.  Now,  in  order 


158  INORGANIC  CHEMISTRY.  [§  101. 

to  measure  the  force  with  which  an  action  tends  to  proceed, 
we  often  make  use  of  an  opposing  force  of  known  magnitude, 
which  is  just  great  enough  to  stop  the  action.  If  this  opposing 
force  is  too  small,  the  internal  driving  force  of  the  system  will 
overpower  it  and  thereby  do  a  certain  amount  of  work,  and  this 
amount  of  work  will  be  the  greater,  the  greater  the  counter  force 
that  is  overcome,  or  in  other  words,  the  smaller  the  difference 
between  this  counter  force  and  the  driving  force  of  the  system. 
For  measuring  affinity  we  can  thus  make  use  of  the  simple 
mechanical  notions  which  serve  for  the  measurement  of  forces 
in  general,  as,  for  instance,  in  an  ordinary  weighing.  We  oppose 
the  force  to  be  measured  with  another  of  known  but  variable 
magnitude  and  allow  the  latter  to  change  until  equilibrium  is 
established.  There  is  then  equality  between  the  known  force 
and  the  force  to  be  measured. 

The  free  energy  is  in  general  not  equal  to  the  total  energy 
that  comes  into  play  in  a  reaction;  but  frequently  the  difference 
is  not  great,  as,  for  instance,  in  reactions  between  solid  com- 
pounds or  in  solution.  Herein  lies  the  explanation  of  the  excep- 
tions to  BERTHELOT'S  principle  as  well  as  the  reason  for  its 
agreement  with  experiment. 

The  total  energy-content  of  a  body  consists,  according  to  HELM- 
HOLTZ,  of  free  and  bound  energy.  The  free  energy  alone  is  capable 
of  transformation  into  other  forms  of  work.  The  bound  energy  is 
involved  in  such  changes  as  those  of  state.  When  ice  melts  a  con- 
siderable amount  of  heat  is  absorbed  which  cannot  be  transformed 
into  work,  but  only  seems  to  increase  the  molecular  movements  of  the 
water  molecules.  The  bound  energy  of  water  is  therefore  greater  than 
that  of  ice  at  the  same  temperature.  Similarly,  there  are  various  other 
processes  where  the  bound  energy  is  changed. 

It  can  be  proved  theoretically  that  in  every  action  proceeding  of  its 
own  accord  the  free  energy  must  decrease. 

In  the  case  of  an  exothermic  reaction  the  evolution  of  heat  is  due 
in.  part  to  the  decrease  of  the  free  energy  of  the  system.  Further,  the 
bound  energy  can  at  the  same  time  either  be  partly  converted  into 
heat,  remain  unchanged,  or  increase  less  than  the  decrease  of  free 
energy  calls  for;  however,  if  the  decrease  of  free  energy  in  a  reac- 
tion is  less  than  the  increase  of  bound  energy,  the  whole  caloric 
effect  must  be  negative,  which  is  to  say,  that  the  reaction  is  endo- 
thermic. 


§  101.]  THERMOCHEMISTRY.  159 

For  the  measurement  of  affinity  it  is  therefore  necessary  to 
determine  this  maximum  work  or  free  energy  which  is  involved 
in  chemical  reactions.  Two  means  are  available,  one  the  deter- 
mination of  the  electromotive  force  that  can  be  created  by  it, 
and,  secondly,  the  determination  of  the  equilibrium  constant  of 
the  reaction  in  question. 

We  shall  learn  in  the  chapter  on  electrochemistry  that 
reactions  can  in  many  cases  be  conducted  so  as  to  produce  an 
electric  current.  If  the  reaction  is  reversible,  it  can  be  brought 
to  a  stop  by  sending  a  current  of  the  same  energy  through  the 
system  in  the  opposite  direction.  The  energy  of  an  electric 
current  is  represented  by  the  product  of  two  factors,  the  amount 
of  electricity  (expressed  in  coulombs)  and  the  electromotive 
force  (expressed  in  volts).  Now  the  decomposition  of  an  equiv- 
alent amount  of  each  compound  requires,  according  to  FARADAY'S 
law,  the  same  amount  of  electricity,  namely,  96,540  coulombs 
per  equivalent  weight;  whence  it  follows  that  the  electromotive 
force  must  be  proportional  to  the  affinity ;  in  other  words,  that 
the  electromotive  force  is  a  measure  of  the  affinity.  Accordingly, 
the  affinity  which  seeks  to  bring  about  a  chemical  transformation 
must  be  opposed  by  an  electromotive  force  just  great  enough 
to  prevent  the  reaction.  This  electromotive  force  is  then  the 
exact  measure  of  the  affinity  whose  action  it  prevents. 

The  free  energy  or  maximum  amount  of  work  which  the 
reaction  produces  is  accordingly  equal  to  the  energy  of  the 
electric  current  produced. 

The  second  general  method  for  measuring  affinity  is  applicable 
in  all  cases  involving  a  chemical  equilibrium.  We  learn  from 
thermodynamics  that  the  equilibrium  constant  K  and  the  maxi- 
mum amount  of  work  A  done  by  the  reaction,  bear  the  following 
relation  to  each  other: 

A  =  RTlogeK, 

when  unit  concentrations  of  the  reacting  substances  are  in- 
volved. R  is  the  gas  constant  (§  35)  and  T  the  absolute 
temperature.  K  is  also  dependent  on  (i.e.,  a  function  of)  the 
temperature. 

The  student  will  find  it  interesting  to  learn  from  the  ap- 
propriate text-books  of  physical  chemistry  how  these  two 


160  INORGANIC   CHEMISTRY.  [§§  102- 

methods  are  utilized  for  the  ]  calculation  of  affinity  in  a  variety 
of  special  cases. 


THE  DISPLACEMENT  OF  EQUILIBRIUM. 

1 02.  When  two  systems  are  in  equilibrium  with  each  other 
(e.g.,  2H2  +  O2<=±2H2O),  the  position  of  this  equilibrium  is  depend- 
ent on  various  circumstances.     The  relationship  is  expressed  by 
the  rule  of  LE  CHATELIER: 

When  any  system  is  in  physical  or  chemical  equilibrium,  a 
change  in  one  of  its  equilibrium  factors  produces  a  change  in  the 
system,  whose  effect  is  opposite  to  that  of  the  former  change. 

This  rule,  or  theorem,  which  can  be  called  the  principle  of  the 
resistance  of  the  reaction  to  the  action,  furnishes  us  with  a  con- 
venient means  of  foretelling  in  many  instances  the  direction  which 
a  reaction  will  follow.  Some  examples  may  be  given  to  illustrate 
the  rule. 

(1)  When  a  system  of  water  and  ice  is  subjected  to  increased 
pressure,  the  ice  melts;  that  is,  that  process  goes  on  which  involves 
a  contraction,  for  by  this  contraction  the  system  diminishes  the 
pressure  exerted  on  it. 

(2)  Monoclinic  sulphur,  when  compressed  near  the  transition 
point  (the  temperature  of  equilibrium  for  ordinary  pressure) ,  passes 
over  into  rhombic  sulphur,  since  this  process  involves  a  lessening 
of  volume,  and  in  the  end  also  a  diminution  of  pressure,  as  in  the 
previous  case. 

(3)  When  a  solution  is  diluted  the  osmotic  pressure  decreases 
according  to  BOYLE'S  law;  in  the  case  of  a  solution  of  an  electro- 
lyte dilution  will  be  followed  by  further  dissociation,  since  this 
increases  the  osmotic  pressure. 

(4)  When  a  liquid  is  heated,  more  vapor  is  formed;  since  the 
vaporization  absorbs  heat    its  effect  is  opposite  to  that  of  the 
heating. 

(5)  In  partially  dissociated  N204  an  increase  of  pressure  drives 
back  the  dissociation,  while  diminution  of  pressure  increases  the 
dissociation.     The    former    change    carries    with    it    a    pressure 
decrease,  the  latter  a  pressure  increase. 

103.  VAN'T  HOFF'S  principle  of  mobile  equilibrium  is  a  special 
case  of  LE  CHATELIER'S  rule,  but  was  derived  from  thermodynamics 


104.]  THERMOCHEMISTRY.  161 

independently.     It   says:    An   equilibrium   between   two   different 
states  of  matter  (systems)  displaces  itself  under  constant  pressure  by 

a  J—r-  of  temperature  to  that  one  of  the  two  systems  whose  formation 
rise 

evolves  T 

_T r  heai.     A  few  examples  will  serve  to  make  this  clear 

absorbs 

(1)  Rhombic  sulphur  becomes  monoclinic  when  heated  above 
the  transition  point,  since  heat  is  absorbed  by  this  transition. 
Below    this    temperature    the    inverse    transition    takes  place. 
(Ordinary  pressure   is   assumed   in   each  case.)      The   reaction 
works  in  opposition  to  the  temperature  change  produced  from 
without. 

(2)  A  salt  whose  heat  of  solution  is  negative  (saltpetre)  dis- 
solves to  a  greater  degree  if  the  temperature  rises.     If  its  heat  of 
solution  is  positive,  a  rise  of  temperature  causes  a  separation  from 
solution.    (§  235). 

This  principle  leads  to  a  further  very  remarkable  deduction. 
Since  an  elevation  of  the  temperature  requires  the  displacement 
of  the  equilibrium  in  the  direction  of  that  system  which  is  formed 
with  absorption  of  heat,  endothermic  reactions  must  predominate  at 
high  temperatures.  On  the  other  hand,  exothermic  reactions  must 
be  generally  associated  with  low  temperatures.  From  the  mathe- 
matical formulation  of  the  principle  it  follows  that  at  the  absolute 
zero  all  reactions  must  be  exothermic.  At  the  prevailing  room  tem- 
perature, which  is  not  so  very  far — approximately  300° — above 
the  absolute  zero,  most  reactions  are  still  exothermic,  although 
endothermic  reactions  do  occur  (formation  of  C12O,  etc.).  At 
the  temperature  of  the  electric  arc  (2000-2500°-)  exothermic 
compounds  are  mostly  incapable  of  existence,  endothermic  com- 
pounds being  obtained.  We  saw  above  that  ozone  and  hydrogen 
peroxide,  both  endothermic,  are  formed  at  a  very  high  tempera- 
ture (§§  36  and  38);  and  we  shall  later  see  that  the  endothermic 
compounds  nitric  oxide  (§§  120  and  127)  and  acetylene  (§  180) 
can  be  synthesized  in  the  heat  of  the  electric  arc. 

PASSIVE  RESISTANCES.. 

104.  A  reaction  can  only  proceed  of  itself  in  case  it  yields  free 
energy.  It  could  no  more  do  otherwise  than  a  stone  could  fly 
up  into  the  air.  Still  we  observe  that  certain  reactions  which 


162  INORGANIC  CHEMISTRY.  [§§  104- 

would  undoubtedly  liberate  free  energy  do  not  occur.  Accord- 
ingly, we  have  to  assume  that  circumstances  can  arise  to  prevent 
the  occurrence  of  a  strictly  possible  reaction.  Such  hindering 
conditions  can  be  comprehensively  termed  passive  resistances. 
Their  effect  is  noticed,  for  instance,  in  the  retarding  of  reaction 
velocity,  especially  at  low  temperatures.  We  saw  an  example  of 
this  in  §  12  in  a  mixture  of  hydrogen  and  oxygen.  Moreover 
PICTET  has  shown  that  sodium,  which  reacts  rather  vigorously 
with  alcohol  at  ordinary  temperatures,  floats  on  it  quietly  at 
—80°  without  any  apparent  reaction;  even  concentrated  hydro- 
chloric acid  and  marble  do  not  react  upon  each  other,  or  at 
least  only  very  slowly,  when  they  are  cooled  to  a  low  tempera- 
ture. In  general,  the  reaction  velocity  lessens  as  the  temperature 
falls.  It  is  to  this  circumstance  in  many  cases  that  we  have  to 
attribute  the  non-occurrence  of  reactions  which  are  thermodynam- 
ically  possible. 

It  has  been  found  that  the  variation  of  the  reaction  velocity 
with  the  temperature  may  in  general  be  expressed  thus:  when  the 
temperature  increases  arithmetically,  the  velocity  increases  geo- 
metrically. Experience  has  also  shown  that  a  temperature  rise  of 
ten  degrees  in  the  neighborhood  of  room  temperature  generally 
involves  about  a  doubling  or  trebling  of  the  reaction  velocity. 

It  is  not  difficult  to  see  what  an  exceedingly  important  role  the 
passive  resistances  above  referred  to  play  in  nature.  Were  it  not 
for  them,  the  phenomenon  of  combustion  and  the  oxidation  of 
metals,  etc.,  could  take  place  at  ordinary  temperatures;  every- 
thing combustible  would  then  burn  and  there  could  be  no  animal 
and  vegetable  life  on  the  earth. 

NITROGEN. 

105.  This  element  occurs  free  in  the  air,  which  contains  about 
80%  nitrogen  and  20%  oxygen.  In  combination,  it  is  found  in  the 
salts  of  nitric  acid,  e.g.  saltpetre,  and  also  in  the  albuminoids, 
which  form  an  important  constitutent  of  animal  and  vegetable 
organisms. 

Nitrogen  can  be  easily  isolated  from  the  air  by  removing  the 
oxygen.  This  is  accomplished  in  various  ways.  Phosphorus, 
when  burned  in  the  air,  absorbs  the  oxygen  to  form  phosphorus 


106.]  NITROGEN.  163 

pentoxide,  and  the  residual  gas,  aside  from  slight  admixtures 
(§  110),  is  nitrogen.  Again,  air  can  be  passed  over  heated  copper 
in  a  finely  divided  condition,  whereupon  copper  oxide  is  formed  and 
nitrogen  left. 

In  this  process  the  oxygen  of  the  air  soon  converts  all  the  copper  into 
copper  oxide,  so  that  of  course  only  a  limited  amount  of  nitrogen  can  thus 
be  obtained  with  the  aid  of  a  given  amount  of  copper.  However,  it 
the  air  is  first  passed  through  ammonia  water,  the  process  can  be  carried 
on  continuously,  since  the  hydrogen  of  the  ammonia,  NH3,  constantly 
reduces  the  oxidized  copper. 

Copper  can  also  absorb  the  oxygen  of  the  air  at  ordinary  tem- 
peratures, if  it  is  treated  with  a  solution  of  ammonia  and  am- 
monium carbonate.  Moist  phosphorus  combines  with  oxygen  even 
at  ordinary  temperatures,  so  that  a  volume  of  air  which  remains 
in  contact  with  pieces  of  phosphorus  for  some  minutes  loses  its 
oxygen.  An  alkaline  solution  of  pyrogallol  also  has  the  ability  to 
absorb  oxygen  at  ordinary  temperatures.  These  reactions  are 
made  use  of  in  gas  analysis. 

1 06.  Pure  nitrogen  is  obtained  by  the  direct  decomposition  of 
certain  of  its  compounds,  especially  by  heating  ammonium  nitrite. 

NH4NO2  =  N2  +  2H2O. 

This  is  usually  accomplished  by  boiling  a  solution  of  equal  parts  by 
weight  of  potassium  nitrite,  KN02,  sal  ammoniac,  NH4C1,  and  potassium 
dichromate,  K2Cr207,  in  3  parts  of  water.  The  NH4C1  and  KNO2  react  to 
form  KC1  and  NH4N02. 

By  heating  ammonium  chromate,  (NH4)2Cr04  (a  mixture  of 
ammonium  chloride  and  potassium  dichromate  is  more  convenient), 
nitrogen  is  also  set  free: 

K2Cr  207 + 2NH4C1 =  N2  +  Cr203 + 2KC1 + 4H2O. 

An  example  of  the  formation  of  nitrogen  by  the  indirect  decom- 
position of  its  compounds  is  the  reduction  of  nitrogen  oxides 
by  hot  copper: 

2NO-f2Cu=N2+2CuO. 


164  INORGANIC  CHEMISTRY.  [§§106- 

Physical  Properties. — Nitrogen  is  a  colorless  and  tasteless  gas. 
Its  specific  gravity  based  on  air  is  0.9682,  its  density  compared 
with  hydrogen  is  therefore  13.X93.  1  1.  N  weighs  1.2521  g.  at  0° 
and  760  mm.  It  is  one  of  the  most  difficult  gases  to  condense, 
its  critical  temperature  being  —146°.  Its  boiling-point  is  —194°. 
At  —214°  it  becomes  solid.  It  is  only  slightly  soluble  in  water, 
even  less  so  than  oxygen. 

Chemical  Properties. — Nitrogen  is  chemically  very  indifferent; 
it  unites  with  no  element  at  ordinary  temperatures  and  at  higher 
temperatures  with  only  a  few.  Boron,  silicon,  titanium,  barium, 
strontium,  calcium,  magnesium,  chromium  and  also  certain  rare 
elements  combine  directly  with  nitrogen  at  red  heat,  forming 
nitrides.  The  direct  union  of  nitrogen  with  a  large  number 
of  elements,  notably  metals,  is  best  accomplished  by  a  special 
experimental  arrangement,  as  follows:  In  a  liquid  mixture  of  90% 
argon  and  10%  nitrogen  an  electric  arc  is  produced  between  a  silver 
anode  and  the  element  concerned.  The  low  temperature  of  the 
bath  prevents  the  decomposition  of  the  nitrides  formed  in  the  arc. 

Nitrogen  unites  with  oxygen  under  the  influence  of  induc- 
tion sparks  directly  (reddish  brown  N02  being  formed);  with 
hydrogen  it  combines  in  a  similar  way.  When  a  mixture  of 
hydrogen  and  nitrogen,  together  with  a  few  drops  of  concen- 
trated hydrochloric  acid,  is  introduced  into  a  tube  over  mercury 
and  induction  sparks  are  sent  through,  clouds  of  ammonium 
chloride,  NH4C1,  are  produced,  the  nitrogen  and  hydrogen  having 
united  to  form  ammonia,  NH3.  These  last  two  reactions  and 
the  fact  that  nitrogen  is  not  able  to  support  combustion  serve  for 
the  identification  of  nitrogen  gas. 

The  molecule  of  nitrogen  consists  of  two  atoms,  this  having 
been  demonstrated  in  the  same  way  as  for  oxygen  and  other  gaseous 
elements. 


107.]  THE  ATMOSPHERE.  165 


THE  ATMOSPHERE. 

107.  The  air  was  regarded  as  an  element  up  to  the  end  of  the 
eighteenth  century.  It  finally  developed  from  the  investigations 
of  PRIESTLEY  and  LAVOISIER  that  it  is  not  a  simple  body.  The 
correct  explanation  of  the  phenomena  of  combustion  led  to  this 
conclusion. 

Before  LAVOISIER'S  time  the  explanation  of  the  phenomena  of  com= 
bustion  was  just  the  reverse  of  the  present  one.  It  was  then  thought 
that  all  combustible  or  oxidizable  substances  had  a  common  constituent, 
phlogiston.  According  to  this  theory,  which  was  presented  by  STAHL 
(1660-1734),  the  combustion  of  a  body  is  due  to  the  escape  of  phlogiston. 
If  this  occurs  in  a  violent  manner  we  have  the  phenomenon  of  fire.  The 
more  inflammable  a  substance  is,  the  more  phlogiston  it  was  supposed  to 
contain.  Sulphur,  phosphorus,  carbon  and  hydrogen  therefore  ranked 
as  being  very  rich  in  phlogiston.  As  to  the  real  nature  of  this  phlogiston 
opinions  were  decidedly  different.  At  various  times  experiments  were 
performed  with  the  hope  of  isolating  the  substance.  For  a  while  it  was 
thought  with  CAVENDISH  that  hydrogen  was  pure  phlogiston. 

The  prevailing  ideas  were  as  follows:  Substances  that  possess  much 
phlogiston  can  transfer  it  to  those  which  have  none  or  very  little.  The 
metals,  for  example,  are  substances  that  contain  a  certain  amount  of 
phlogiston,  which  they  give  off  on  being  heated  in  the  air;  by  this  process 
they  are  changed  to  calxes  (now  called  oxides) ,  which  contain  no  phlogis- 
ton. When  one  of  these  calxes  is  heated  with  carbon  or  hydrogen,  it 
absorbs  phlogiston  from  them  and  is  changed  back  again  to  the  metal. 
The  fact  that  sulphur,  phosphorus  or  any  other  inflammable  substance 
soon  ceases  to  burn  when  it  is  enclosed  in  an  air-tight  space  was  explained 
by  the  supposition  that  the  air  has  then  become  so  saturated  with  phlogis- 
ton that  the  latter  can  no  longer  escape  from  the  burning  body. 

We  see  from  the  above  that  this  theory  led  men  to  view  many  phenom- 
ena from  a  common  standpoint  and  undoubtedly  contributed  in  no  small 
degree  to  the  advancement  of  chemistry.  So  long  as  the  phenomena  of 
burning  were  regarded  in  that  light,  there  was  no  occasion  to  doubt  the 
elemental  nature  of  air.  They  believed  that  bodies  lose  something  when 
they  are  burned,  while  we  now  know  that  on  the  contrary  something  is 
taken  up  from  the  air.  The  great  mistake  of  the  phlogiston  theory  was, 
that  it  did  not  regard  the  increase  in  weight  of  the  burned  body;  as  soon 
as  LAVOISIER  and  others  drew  attention  to  this  most  important  fact,  the 
phlogiston  theory  could  no  longer  be  upheld. 

On  the  first  of  August,  1774,  PRIESTLEY  had  discovered  oxygen,  which 


166  INORGANIC  CHEMISTRY.  [§§  107- 

he  himself  regarded  as  air  devoid  of  phlogiston  ("  dephlogisticated  air  ") ;  * 
LAVOISIER,  however,  recognized  this  substance  as  the  essential  principle 
of  all  burning  and  oxidation.  It  now  required  only  a  step  to  reach  the 
conception  that  air  is  not  an  element,  but  contains  another  gas  in  addition 
to  oxygen,  and  that  this  gas  does  not  support  combustion.  The  experi- 
ment by  which  LAVOISIER  demonstrated  this  has  been  already  described 
(§  9).  By  measuring  the  amount  of  nitrogen  which  remained  after  the 
absorption  of  the  oxygen  by  hot  mercury  he  was  able  to  determine  fairly 
accurately  the  composition  of  air. 

1 08.  Constituents  of  the  Atmosphere. — Besides  oxygen  and 
nitrogen  air  contains  argon  and  the  other  elements  described  in 
§  110,  hydrogen,  and  also  variable  amounts  of  water  vapor, 
carbon  dioxide  (very  nearly  0.04%  on  the  average),  ammonia, 
ozone,  and  perhaps  hydrogen  peroxide  (the  last  three  in  extremely 
small  quantities).  Incidentally  sulphur  dioxide  and  other  gases 
are  found  in  the  air  (e.g.,  in  the  vicinity  of  volcanoes).  The 
lower  strata  of  air  always  contain  floating  dust  particles,  microbes, 
etc. 

Analysis  of  Air. — The  proportional  amounts  of  oxygen  and 
nitrogen  in  carefully  dried  air,  free  from  carbon  dioxide,  etc.,  have 
been  repeatedly  determined  with  all  due  precaution.  According 
to  the  method  of  DUMAS  and  BOUSSINGAULT  this  can  be  done 
as  follows : 

The  tube,  ab  (Fig.  34) ,  containing  copper  turnings  is  connected 
with  the  globe,  Vt  all  air  having  been  removed  from  both.  The  end 
of  the  tube  marked  6  is  attached  to  the  various  pieces  of  apparatus 
Cj  B  and  A,  which  are  to  remove  the  carbon  dioxide  and  water 
vapor  from  the  inflowing  air.  The  globe,  V,  is  first  carefully 
weighed  without  air.  Thereupon  the  tube  is  heated  by  means  of  a 
furnace  and  a  slow  current  of  air  is  allowed  to  pass  through  it  to 
the  globe  by  partially  opening  the  stop-cocks  u  and  r,  the  oxygen 
being  meanwhile  asborbed  by  the  hot  copper.  By  subsequently 
weighing  the  globe  the  amount  of  nitrogen  which  it  contains  can 
be  determined  and  by  weighing  the  tube  before  and  after  we  can 

*  From  the  letters  and  laboratory  notes  of  SCHEELE,  published  by  Baron 
NORDENSKIOLD  (Stockholm,  1892),  it  is  evident  that  oxygen  was  known  to 
SCHEELE  sooner  than  to  PRIESTLEY;  he  called  it  " Feuerluft."  However,  this 
discovery  did  not  seem  to  lead  him  any  nearer  than  PRIESTLEY  to  a  correct 
understanding  of  the  phenomena  of  burning. 


109-] 


THE  ATMOSPHERE. 


167 


find  the  amount  of  oxygen.  '  In  this  way  the  ratio  of  oxygen  to 
nitrogen  in  air  can  be  ascertained. 


FIG.  34. — ANALYSIS  OF  Am. 

Another  method  is  the  eudiometric  method.  A  known  volume 
of  air  is  mixed  with  a  sufficient  known  volume  of  pure  (electro- 
lytic) hydrogen.  On  allowing  an  electric  spark  to  pass  through, 
the  hydrogen  and  oxygen  unite  to  form  water,  which  is  deposited 
on  the  sides  of  the  vessel.  Inasmuch  as  2  vols.  hydrogen  combine 
with  1  vol.  oxygen,  one-third  of  the  volume  that  disappeared  must 
have  been  oxygen. 

109.  These  and  other  methods  of  investigation  have  shown  that 
the  composition  of  the  air  is  nearly  constant.  In  all  parts  of  the 
earth,  as  well  as  at  the  highest  altitudes  which  balloons  have 
reached,  it  consists  of 

20.81%  oxygen  and  79.19%  nitrogen  by  volume;  and 
23.01%       "         ""    76.99%       "        "  weight. 

The  observed  variations  from  this  ratio  amount  to  hardly  ±0.1%. 
Moreover,  the  composition  does  not  appear  to  change  with  time; 
our  present  analyses  agree  with  those  of  DTJMAS  and  BOUSSINGAULT 
made  in  1841. 

This  result  seems  surprising  at  first  thought,  because  oxygen 
and  nitrogen  are  constantly  being  removed  from  the  air  and  again 
returned  to  it  and  it  does  not  necessarily  follow,  indeed  it  is  rather 
an  improbability,  that  the  losses  and  gains  will  exactly  balance. 


168  INORGANIC    CHEMISTRY.  [§  109 

The  oxygen  passes  through  the  following  cycle:  Free  oxygen 
is  consumed  in  all  sorts  of  oxidations  of  which  the  mineralization 
of  organic  matter  is  the  most  important.  By  the  term  "  minerali- 
zation "  is  meant  the  oxidation  of  the  residues  of  plants  and 
animals  by  the  oxygen  of  the  air  with  the  aid  of  bacilli.  The 
carbon  of  these  residues  is  oxidized  to  carbon  dioxide;  the  nitro- 
gen, phosphorus,  sulphur  and  other  elements  return  to  the 
" mineral"  state,  as  nitrates,  sulphates,  etc.  Along  with  this 
process  there  are  the  other  oxygen-consuming  processes  of  the 
respiration  of  animals  and  plants  and  the  burning  of  fuels,  carbon 
dioxide  being  formed  in  all  cases.  This  carbon  dioxide  is  em- 
ployed by  the  plants  in  their  process  of  assimilation,  the  oxygen 
in  it  being  again  given  back  to  the  air.  It  will  therefore  depend 
on  the  relative  magnitude  of  this  process  as  to  whether  just  as 
much  oxygen  gets  back  into  the  air  as  was  previously  taken  up  in 
the  formation  of  carbon  dioxide.  The  oxygen  which  serves  for 
other  oxidations  does  not  necessarily  return  to  the  air. 

Different  investigators  have  attempted  to  estimate  the  amount  of 
carbon  which  annually  enters  into  the  cycle  of  organic  life.  DUBOIS 
calculated  that  every  year  the  plants  assimilate  118.5  million  million 
kilograms  C02,  which  is  almost  TO  of  the  total  carbon  dioxide  in  the 
atmosphere.  The  amount  of  C02  given  off  by  the  entire  animal  world  is 
estimated  at  2.5  million  million  kilograms,  which  brings  us  to  the 
startling  result  that  only  about  2%  of  the  existing  plant  material  is 
engaged  in  the  cycle  with  the  animal  life.  All  the  rest  of  the  carbon 
dioxide  required  by  the  plants  comes  from  the  process  of  mineralization. 
The  amount  of  carbon  dioxide  produced  by  the  burning  of  coal,  etc.,  is 
estimated  by  CREDNER  at  1.3  million  million  kilograms. 

Nitrogen  passes  through  a  cycle  too.  Most  of  the  nitrogen 
that  occurs  in  the  form  of  organic  compounds  in  animal  and  vege- 
table tissues  remains  in  the  combined  state  after  the  death  of 
the  organism,  either  as  ammonia  or  as  nitric  acid  or  in  other  nitro- 
genous products.  During  the  process  of  decay  the  combined 
nitrogen  is  partially  liberated;  in  the  burning  of  plant  and  animal 
remains  all  of  it  is  set  free.  On  the  other  hand,  certain  plants,  the 
Leguminosce,  are  able  by  symbiosis  with  bacteria  to  absorb  free 
nitrogen  from  the  air  directly.  There  are  also  bacteria  which, 
acting  alone,  can  assimilate  nitrogen.  Moreover,  in  storms  some 


§  109-]  THE  ATMOSPHERE.  .  169 

nitrogen  combines  with  oxygen,  and  again,  silent  electric  discharges, 
such  as  must  frequently  pass  between  earth  and  clouds,  cause  the 
nitrogen  to  enter  into  combination.  Here  the  question  again 
arises  whether  as  much  comes  back  to  the  air  as  goes  out. 

From  what  has  been  said  it  is  sufficiently  clear  that  it  would 
fos  a  mere  coincidence  if  exactly  as  much  oxygen  should  happen 
to  be  withdrawn  as  is  given  back.  Approximate  compensation 
p: :  oably  takes  place,  but,  even  if  it  should  not,  the  atmosphere  is 
PC  vast  that  its  composition  would  be  only  slightly  affected  in  the 
course  of  centuries. 

The  following  calculation  will  convince  one  of  the  soundness  of  this 
argument:  The  normal  atmospheric  pressure  is  760  mm.  mercury;  this 
is  due  to  the  weight  of  the  air  and  the  moisture  in  it.  Granted  that  the 
pressure  of  the  latter  averages  10  mm.,  we  have  750  mm.  left  for  the 
pressure  of  the  air  itself;  i.e.  the  weight  of  the  air  is  equal  to  that  of  a 
layer  of  mercury  750  mm.  thick  extending  over  the  entire  surface  of  the 
earth.  This  weight  can  be  calculated  thus:  The  volume  of  the  space 
between  two  concentric  spheres  is  4;r.R2r,  if  R  is  the  radius  of  the  inner 
sphere,  and  r  the  thickness  of  that  space.  The  radius  of  the  earth  (R) 
is,  on  the  average,  6,370,284m.;  r  is  0.75m.;  therefore,  taking  into 
consideration  the  specific  gravity  of  mercury  (13.59),  we  have  for  the 
desired  weight  of  mercury  or  air  5.2  XlO18  kilograms.  Since  1  m.* 
air  at  0°  and  760  mm.  pressure  weighs  1.2932  kg.,  the  above  weight 
corresponds  to  a  volume  of  air  of  4 XlO18  m.3  (at  0°  and  760  mm.) 
or  f  XlO18  =8 X 1017  m.3  of  oxygen.  In  comparison  with  this  the  amount 
of  oxygen  which  is  withdrawn  from  the  air  in  breathing,  burning,  etc.,  is 
very  small,  as  may  be  seen  from  the  figures  on  the  preceding  page  for 
the  quantities  used  by  animals  and  plants.  Since,  on  the  other  hand, 
the  assimilative  process  of  the  plants  yields  a  considerable  amount  in 
addition,  ths  variations  in  the  proportion  of  oxygen  in  the  air  must 
obviously  be  imperceptible  with  our  present  analytical  methods. 

The  air  is  a  mixture.  It  cannot  be  a  compound  of  nitro- 
gen and  oxygen  for  the  following  reasons:  (1)  the  ratio  of  nitrogen 
to  oxygen  is  different  than  it  would  be  for  a  compound  of  the 
two  elements,  for  in  the  latter  case  it  would  have  to  correspond 
to  the  ratio  of  the  atomic  weights  or  a  multiple  of  the  same;  (2) 
by  mixing  nitrogen  and  oxygen  in  xthe  ratio  in  which  they  exist  in 
air  a  synthetical  air  is  obtained  which  is  in  every  respect  like  that 
around  us.  (This  excludes  the  possibility  of  air  containing  a  per- 
ceptible amount  of  a  compound  of  the  two  elements  in  addition  to 


170  INORGANIC  CHEMISTRY.  [§§  109- 

free  nitrogen  and  free  oxygen.)  (3)  The  ratio  of  the  solubilities 
of  the  oxygen  and  the  nitrogan  of  the  air  in  liquids  is  the  same 
as  that  calculated  from  the  solubilities  of  the  pure  gases  oxygen 
and  nitrogen,  after  taking  into  account  their  partial  pressures. 
This  could  not  be  the  case  if  the  air  contained  a  compound  of 
oxygen  and  nitrogen;  (4)  when  liquid  air  boils  the  first  part  of 
the  distillate  is  chiefly  nitrogen. 

The  liquefaction  of  air  is  now  carried  on  in  commerce.  The 
methods  used  by  LINDE  and  by  HAMPSON  are  based  on  the  same 
principle,  namely,  cooling  the  air  by  expansion.  Further  details 
may  be  found  in  text-books  on  physics. 

Liquid  air  is  very  mobile  and  has  a  bluish  tint.  It  is  usually 
somewhat  cloudy  because  of  suspended  particles  of  ice  (congealed 
atmospheric  moisture)  and  solid  carbon  dioxide.  These  may  be 
removed  by  filtration  through  filter-paper.  It  boils  at  about 
— 190°.  It  is  now  extensively  used  in  producing,  and  demonstrat- 
ing the  effects  of,  very  low  temperatures.  When  carbon  dioxide, 
for  example,  is  led  into  a  flask  containing  liquid  air,  it  falls  in 
the  solid  form  like  snow-flakes.  In  spite  of  its  low  temperature 
liquid  air  can  be  poured  upon  the  hand  without  danger;  it  does 
not  even  feel  cold  (on  account  of  the  LEYDENFROST  phenomenon). 
Liquid  air  is  much  richer  in  oxygen  than  the  gaseous  air  of  the 
atmosphere,  containing  about  50%.  If  a  glowing  splinter  is 
dipped  into  the  liquid,  the  wood  begins  to  burn  very  vigorously 9 
producing  a  violent  reaction.  It  can  be  preserved  for  a  rather 
long  time  in  vacuum  flasks.  By  fractional  distillation  of  liquid 
air  practically  pure  oxygen  and  nitrogen  can  be  obtained. 
According  to  ERDMANN  pure  nitrogen  is  obtained  in  the  cooling 
down  of  liquid  air,  whereupon  nitrogen  crystallizes  out. 

ARGON,  HELIUM  AND  COMPANION  ELEMENTS. 

no.  Argon.  Despite  the  fact  that  air  had  been  already 
analyzed  times  without  number,  it  was  first  discovered  in  the 
course  of  investigations  by  RAYLEIGH  and  RAMSAY  in  1894  that 
there  are  other  elements  in  the  air  than  nitrogen  and  oxygen. 
One  of  these,  named  argon  by  its  discoverers,  is  even  found  to 
the  extent  of  0.9%  by  volume,  or  1.2%  by  weight.  It  was 
on  account  of  its  extraordinary  resemblance  to  nitrogen  that  it 


110.]        ARGON,  HELIUM  AND  COMPANION  ELEMENTS.       171 

was  so  long  overlooked.  The  first  indication  of  its  presence  was 
the  observation  that  the  specific  gravity  of  the  nitrogen  isolated 
from  the  air  is  somewhat  higher  than  that  of  the  nitrogen  pre- 
pared from  ammonium  nitrite  and  other  compounds.  1  liter  of 
nitrogen  from  air  weighed  1.2572  g.,  while  the  same  amount  from 
chemical  compounds  weighed  1.2521  g.;  in  both  cases  at  0°  and 
760  mm.  There  must  therefore  be  another  gas  heavier  than 
nitrogen,  mixed  in  with  the  nitrogen  of  the  air. 

One  of  the  simplest  methods  for  obtaining  argon  from  the 
air  is  to  heat  air  with  a  mixture,  of  1  g.  magnesium,  0.25  g. 
sodium  and  5  gr.  freshly  ignited  lime.  On  account  of  the  high 
temperature  free  calcium  is  formed  : 

Mg+CaO  =  MgO  +  Ca, 

and  it  is  in  such  a  finely  divided  condition  that  it  absorbs  oxygen 
greedily  and  also  nitrogen,  so  that  only  argon  is  left. 

Argon  can  also  be  isolated  with  the  help  of  calcium  carbide. 
When  calcium  carbide  (better,  mixed  with  10%  calcium  chloride) 
is  brought  in  contact  with  air  at  about  800°,  it  absorbs  both 
oxygen  and  nitrogen: 


2CaC2  +  O2  =  2CaO  +  4C  ;  CaC2  +  N2  =  CaCN2  +  G. 

This  is  a  suitable  method  for  preparing  argon  in  large  quan- 
tities. 

After  argon  had  been  once  discovered  it  was  found  elsewhere 
than  in  the  atmosphere  ;  some  mineral  waters  contain  it  in  solution, 
certain  rare  minerals  yield  it  when  heated,  etc. 

Argon  is  a  colorless,  odorless  gas,  having  a  vapor  density  of 
19.957.  It  has  been  condensed  to  a  colorless  liquid,  that  boils  at 
—  186.9°,  by  cooling  with  boiling  oxygen  and  compressing  to  about 
50.6  atmospheres;  it  solidifies  at  —189.6°.  It  is  somewhat  more 
soluble  in  water  than  is  nitrogen  (0.05780  parts  in  1  vol.  at  0° 
and  760  mm.  pressure).  As  to  its  chemical  nature,  it  is  interest- 
ing that  no  one  has  yet  succeeded  in  preparing  a  compound  of  argon. 

It  is  certain  that  what  is  now  called  argon  is  neither  a  mixture 
nor  a  compound,  but  an  element.  The  boiling-point  and  the1 
melting-point  are  constant,  and  the  vapor  pressure  of  argon  like- 


172  INORGANIC   CHEMISTRY.  [§§110- 

wise  remains  constant  during  the  liquefaction,  so  long  as  any  gas 
is  present.  Moreover,  when  a  certain  volume  of  argon  is  three- 
fourths  dissolved  in  water,  the  undissolved  gas  shows  exactly  the 
same  spectrum  as  the  dissolved.  All  of  the  above  are  charac- 
teristics of  a  homogeneous  substance.  The  extraordinary  stability 
of  the  gas  in  the  presence  of  all  sorts  of  reagents  is  a  strong 
argument  against  its  being  a  compound. 

in.  After  the  discovery  of  argon  RAMSAY  and  TRAVERS  detected  four 
other  rare  gases  in  the  atmosphere,  though  their  quantity  is  very  small. 
These  are  helium,  neon,  krypton,  and  xenon.  In  a  spectroscopic  inves- 
tigation (§  265)  NORMAN  LOCKYER  had  detected  in  the  atmospheres  of 
the  sun  and  many  fixed  stars  considerable  quantities  of  a  gas  unknown  on 
the  earth ;  he  named  it  helium.  In  1895  RAMSAY  and  TRAVERS  succeeded, 
however,  in  obtaining  it  in  small  amounts  on  heating  the  rare  mineral 
cleveite.  Afterward  it  was  also  met  with  as  a  companion  of  argon  in 
certain  other,  chiefly  uraniferous,  minerals  as  well  as  in  mineral  springs, 
for  instance,  those  of  Bath;  and  at  last  it  was  also  discovered  in  the  air. 
At  ordinary  temperatures  helium  is  a  colorless  gas.  It  is  of  all  gases  the 
most  difficult  to  condense;  yet  KAMERLINGH  ONNES  recently  achieved 
the  task.  Helium  boils  at  ^4°  absolute  temperature  (—269°  C.).  By 
quickly  evaporating  it  a  temperature  of  ^  1 .5°  absolute  was  reached, 
which  is  the  lowest  thus  far  attained.  The  critical  pressure  is  ^  3  atmos- 
pheres. Dens,  at  4.29°  abs.  =0.122.  In  water  helium  is  less  soluble 
than  argon.  For  its  relation  to  radium,  see  §  267. 

Helium  and  neon  (0.00086  wt.  %  of  air)  are  found  in  the  most  volatile 
part  of  liquid  air.  DEWAR  proved  that  helium  and  neon  can  be  isolated 
directly  from  the  air  by  bringing  the  air  in  contact  with  ignited  charcoal 
at  —185°.  The  charcoal  has  the  curious  property  of  condensing  in  its 
pores  all  the  other  gases  of  the  air,  and  a  gaseous  residue  is  here  obtained 
which  shows  clearly  the  spectral  lines  of  He  and  Ne. 

While  helium  and  neon  were  found  in  the  most  volatile  part  of  the  air, 
krypton  and  xenon  were  obtained,  on  the  contrary,  from  the 
residue,  after  a  large  quantity  of  liquid  air  had  been  allowed  to  evaporate 
slowly.  Their  separation  was  rendered  possible  by  the  fact  that  krypton 
still  has  a  rather  large  vapor  tension  at  the  temperature  of  liquid  air, 
while  the  vapor  tension  of  xenon  is  then  imperceptible. 

Both  these  elements  occur  only  in  extremely  small  amounts  in  the 
atmosphere.  Krypton  makes  up  0.028%,  xenon  0.005%  (by  weight) 
of  the  air. 

In  the  following  table  some  of  the  data  of  these  elements  are  given. 
The  elements  form  a  natural  group. 


§111.]       ARGON,  HELIUM,  AND  COMPANION  ELEMENTS.       173 


Helium. 

Neon. 

Argon. 

Krypton. 

Xenon. 

Density  (O  =  16)  
Atomic  weight 

1.98 
3  99 

10.1 

20  2 

19.94 

QQ    Qt> 

41.45 

82  9 

65.1 

1  *}O   9 

Boiling-point  at  760  mm.  .  . 

4°  abs. 

86.9°  abs. 

121.  9°  abs. 

163.9°abs. 

These  gases  have  three  properties  in  common  which  are  worthy  of 
mention  here.  In  the  first  place  they  display  characteristic  spectral  lines 
in  PLUCKER  tubes  (§  263),  whereby  it  has  been  possible  to  recognize  them 
and  to  judge  of  their  purity.  In  the  second  place,  no  one  of  these 
elements  has  been  found  to  enter  into  combination  with  other  elements ; 
they  may  therefore  be  considered  nullivalent.  In  the  third  place,  their 
molecule  consists  of  only  one  atom.  This  fact  could  not  be  discovered  in 
the  ordinary  way,  described  in  §  §  33  and  34,  because  of  the  entire  absence 
of  compounds  for  investigation.  It  has,  however,  been  possible  to  ascer- 
tain it  from  the  molecular  heat  of  the  gases.  This  is  the  amount  of  heat 
that  must  be  imparted  to  a  gram  molecule  of  a  gas  in  order  to  raise  its 
temperature  one  degree.  This  quantity  of  heat  differs,  according  as  the 
gas  is  under  constant  pressure  or  under  constant  volume.  It  is  greater 
in  the  first  case  because  under  constant  pressure  the  gas  expands  on 
heating  and  so  does  work,  which  evidently  involves  an  expenditure  of 
heat.  We  saw  in  §  34  that  for  one  gram  molecule  of  a  gas  the  equation 
PV  =  2T  is  applicable,  the  2T  expressing  in' calories  the  external  work 
done  when  a  gas  under  constant  pressure  P  increases  its  volume  by  V,  or 
when  a  gas  being  generated  under  the  pressure  P  comes  to  occupy  a 
volume  V.  If  the  temperature  is  raised  one  degree  we  have  PF  =  2(77  +  1); 
for  each  gram  molecule  of  gas  extra  work  is  therefore  done  equivalent  to 
2  calories.  The  molecular  heat  at  constant  pressure  is  thus  2  cal.  more 
than  that  at  constant  volume. 

From  the  kinetic  theory  of  gases  it  can  be  deduced  that  the  molecular 
heat  of  a  monatomic  gas  at  constant  pressure  is  5  cal.  At  constant 
volume  it  must  be  2  cal.  less,  or  3  cal.  The  ratio  of  these  quantities  of 
heat  is  therefore  5:3  =  1.66.  When  the  molecules  of  the  gas  consist  of 
more  than  one  atom,  more  heat  is  absorbed  for  the  same  rise  of  tem- 
perature, because  heat  is  then  used  not  only  for  the  movement  of  the 
molecules,  but  also  for  that  of  the  atoms  in  the  molecule.  The  ratio  then 
becomes  5+m:3+m,  if  m  is  the  additional  heat.  The  resulting  ratio  is 
thus  less  than  1.66.  By  determining  this  ratio  (which  can  be  found  from 
the  velocity  of  propagation  of  sound  in  the  gas  by  a  well-known  physical 
formula)  we  can  ascertain  whether  the  gases  are  monatomic  or  poly- 
atomic, For  the  gases  of  this  group  the  ratio  was  found  to  be  1.66, 
proving  that  their  molecules  contain  only  one  atom. 


174  INORGANIC  CHEMISTRY.  [§112. 

Compounds  of  Nitrogen  and  Hydrogen. 

ii2.  Until  recent  years  only  one  compound  of  hydrogen  and 
nitrogen  has  been  known,  viz.,  ammonia,  NH3.  At  present,  how- 
ever, we  know  five:  the  others  being  hydrazine,  N2H4,  hydrazoic 
acid,  N3H,  and  the  compounds  of  the  latter  with  ammonia  and  with 
hydrazine  (NH3-N3H  and  N2H4-N3H).  Of  these  five  compounds, 
however,  ammonia  is  by  far  the  most  important. 

AMMONIA. 

The  material  now  used  for  obtaining  ammonia  is  the  "  ammonia 
liquor"  of  the  gas-factories  and  coke  ovens.  The  gases  that 
are  given  off  in  the  dry  distillation  of  coal  are  passed  through 
water,  which  dissolves  the  ammonia. 

In  order  to  obtain  a  pure  ammonia,  the  ammonia  liquor  is  heated  with 
milk  of  lime  and  the  expelled  ammonia  is  led  into  concentrated  sulphuric 
acid.  In  this  way  crystallized  ammonium  sulphate  is  obtained.  It  is 
purified  by  recrystallization  and  again  distilled  with  lime  to  recover  the 
free  ammonia. 

Ammonia  can  be  prepared  synthetically  by  the  following 
methods.  The  direct  synthesis  from  the  elements  was  given 
above  (§  107).  There  are  also  examples  of  its  formation  by  the 
direct  decomposition  of  its  compounds.  Thus  we  obtain  it  by 
heating  the  ammonia  compounds  of  certain  salts,  as  #CaCl2  •  ?/XH3 
and  zAgCl  •  ?/NH3.  A  number  of  organic  compounds  yield  nitrogen 
in  the  form  of  ammonia  on  heating.  Moreover,  ammonia  results 
from  the  action  of  hydrogen  on  certain  nitrogen  compounds,  as, 
for  example,  when  nitric  acid,  HNO3,  comes  in  contact  with 
nascent  hydrogen  (generated  from  zinc  or  iron  filings  and  dilute 
sulphuric  acid),  or  when  a  mixture  of  nitric  oxide,  NO,  with 
hydrogen  is  passed  over  platinum  black: 

2NO  +  5H2  =  2NH3  +  2H20. 

The  formation  of  ammonia  by  the  action  of  free  nitrogen  on 
liydrogen  compounds  has  not  been  brought  about,  but  the  gas  can 
be  produced  by  the  interaction  of  a  hydrogen  compound  with  a 
nitrogen  compound.  An  illustration  of  this  is  the  decomposition 
of  magnesium  nitride  by  water: 

Mg3N2  +  3H20  =  2NH3  +  3MgO. 


§  112..]  AMMONIA.  175 

The  putrefaction  of  organic  matter  (faeces,  urine,  etc.)  evolves 
ammonia.  By  the  action  of  electric .  sparks  on  moist  air  am- 
monium nitrate  is  produced.  These  last  two  methods  of  for- 
mation are  responsible  for  the  slight  traces  of  ammonia  in  the 
air. 

For  the  formation  of  ammonia  from  calcium  cyanamide, 
see  OEG.  CHEM.,  §  266. 

Physical  Properties. — Ammonia  at  ordinary  temperatures  is 
a  gas  with  a  characteristic  odor,  that  excites  one  to  tears.  Its 
specific  gravity  is  8.5  (O=  16)  or  0.589  (air  =  1) ;  1 1.  NH3  at  0°  and 
760  mm.  pressure  weighs  0.76193  g.  It  can  be  easily  liquefied;  it 
boils  at  —33.7°  and  becomes  solid  at  —75°;  it  then  forms  white 
translucent  crystals.  It  is  extremely  soluble  in  water;  at  0°  and 
normal  pressure  1  vol.  H2O  dissolves  1148  vols.,  or  0.875  parts  by 
weight,  of  NH3.  The  specific  gravity  of  the  solution  of  ammonia 
in  water  grows  smaller  as  the  concentration  increases.  The 
evaporation  of  liquid  ammonia  involves  a  considerable  depres- 
sion of  temperature.  This  is  the  principle  of  most  of  the  ice- 
machines  now  in  use. 

Chemical  Properties. — The  characteristic  property  of  this  corn- 
compound  is  that  it  combines  with  acids  directly  to  form  salts: 

NH3 + HC1  =  NH4C1.        NH3  +  HNO3  =  NH4N03. 

Ammonium  Ammonium 

chloride.  nitrate. 

2NH3+H2SO4  =  (NH4)2S04, 

Ammonium 
sulphate. 

In  these  salts  (which  are  almost  all  readily  soluble  in  water) 
the  atomic  group  NH4  plays  the  part  of  a  metal;  they  correspond 
in  every  respect  to  the  compounds  KC1,  KNO3,  K2SO4,  etc.  The 
group,  or  radical,  NH4  has  been  given  a  particular  name;  it  is 
called  ammonium.  More  than  one  attempt  has  been  made  to 
isolate  this  ammonium,  but  always  in  vain.  However,  when 
sodium  amalgam  comes  in  contact  with  a  concentrated  ammo- 
nium chloride  solution,  the  mercury  swells  to  a  soft  spongy  mass 
that  rapidly  decomposes  at  ordinary  temperatures  into  ammonia 
and  hydrogen  and  is  in  all  probability,  therefore,  ammonium  amal- 
gam. If  sodium  amalgam  is  allowed  to  react  with  ammonium 
iodide  dissolved  in  liquid  ammonia  at  —39°,  a  hard  metallic  mass 


176  INORGANIC  CHEMISTRY.  [§§  112- 

is  obtained,  which  swells  with  rising  temperature  because  of  decom- 
position into  mercury,  hydrogen  (1  vol.)  and  ammonia  (2  vols.): 

2NH4=2NH3  +  H2. 

The  aqueous  solution  of  ammonia  reacts  strongly  basic;  so  do 
the  moist  fumes  of  ammonia.  We  must  therefore  assume  that 
this  solution  contains  a  compound  NH4OH,  ammonium  hydroxide, 
and  hence  also  the  ions  NH4  and  OH  in  analogy  with  other  soluble 
bases,  e.g.  potassium  hydroxide,  KOH.  As  a  matter  of  fact, 
however,  the  solution  of  ammonia  conducts  the  electric  current 
much  more  poorly  than  a  solution  of  sodium  hydroxide  of  equiva- 
lent concentration  (§  234).  Ammonium  hydroxide  has  not  yet 
been  isolated.  When  the  solution  of  it  is  evaporated,  NH4OH 
splits  up  into  NH3  and  H20.  Concordant  herewith  is  the  well- 
known  fact  that  ammonia  can  be  entirely  expelled  from  its 
aqueous  solution  by  boiling. 

Ammonia  does  not  burn  in  the  air  but  does  in  oxygen;  in  addi- 
tion to  water  and  nitrogen  traces  of  ammonium  nitrite,  NH4N02, 
and  nitrogen  dioxide,  N02,  are  also  formed.  A  mixture  of 
ammonia  and  oxygen  explodes  violently  when  it  is  ignited.  The 
oxygen  conveyed  by  soil  bacteria  may  also  cause  the  oxidation  of 
ammonia,  producing  nitric  acid.  Chlorine  takes  fire  when  passed 
into  ammonia,  forming  nitrogen,  N2,  and  hydrochloric  acid;  the 
latter  then  unites  with  the  remaining  ammonia  to  form  sal  am- 
moniac, NH4C1.  The  hydrogen  of  ammonia  is  replaceable  by 
metals.  Magnesium,  e.g.  burns  in  ammonia,  forming  magnesium 
nitride,  Mg3N2.  When  ammonia  is  conducted  over  hot  potassium 
or  sodium,  potassium  amide,  NH2K,  or  sodium  amide,  NH2Na,  is 
formed.  These  and  analogous  metal  compounds  are  decomposed 
by  water,  yielding  ammonia  again  and  also  metal  oxide  or 
hydroxide.  At  high  temperatures  (produced  by  induction  sparks) 
ammonia  splits  up  almost  completely  into  its  elements,  the  volume 
being  doubled: 


2  vols.      1  vol.    3  vols. 


On   the   other   hand,    nitrogen   and   hydrogen   can    unite   to   form 
ammonia  under  the  influence  of  induction  sparks  (§  106).     Equilibrium 


114.]  HYDRAZINE,  OR  DIAMIDE.  177 

is  reached  when  3%  of  ammonia  is  formed  in  the  gas  mixture,  N2  +  3H2, 
i.e.,  6%  of  the  theoretical  yield.  This  is  the  reason  why  ammonia  cannot 
be  split  up  by  electric  sparks  to  more  than  97%: 

2NH3^±N2+3H2. 

3  per  cent     97  per  cent 

Nevertheless  N2+3H2  can  be  completely  converted  into  2NH3  by 
induction  sparks  if  the  gas  mixture  is  brought  in  contact  with  an  acid, 
for  by  this  means  ammonia  is  constantly  withdrawn  from  the  gaseous 
system  N2  +  3H2  <=*  2NH3;  the  remaining  gas  mixture  will  therefore  form 
new  NH3  in  order  to  restore  the  equilibrium,  and  so  on,  until  all  the 
nitrogen  and  hydrogen  have  combined. 

Attempts  to  manufacture  ammonia  commercially  by  direct  synthesis 
have  at  last  been  successful.  Some  metal,  such  as  uranium,  iron,  manga- 
nese, or  molybdenum,  is  utilized  as  a  catalyzer  and  the  process  is  carried 
on  at  a  pressure  of  200  atmospheres  and  a  temperature  above  500°.  The 
yield  is  approximately  proportional  to  the  pressure.  Certain  foreign  sub- 
stances are  found  to  act  as  promoters,  others  as  poisons,  of  the  catalysis. 

113.  Composition  of  Ammonia. — If  an  aqueous  ammonia  solu- 
tion (to  which  has  been  added  a  little  sodium  chloride  to  aid 
conduction)  is  subjected  to  electrolysis,  nitrogen,  and  hydrogen 
are  generated  in  the  volume  ratio  of  1:3;  from  this  it  follows 
that  the  molecule  must  contain  3  H-atoms  to  every  1  N-atom, 
i.e.,  the  empirical  formula  is  NH3.  Since  the  specific  gravity 
of  ammonia  gas  is  8.5  (0=16),  the  molecular  weight  is  17,  which 
corresponds  to  the  above  formula. 


HYDRAZINE,  OR  DIAMIDE,  N2H4. 

114.  This  compound  is  now  manufactured  by  the  process  of  RASCHIG. 
He  showed  that  the  reaction  of  sodium  hypochlorite  and  ammonia 
yields  chloramine: 

NH3  -i-NaOCl  =  NH,C1  +  NaOH. 

At  a  low  temperature  in  a  vacuum  this  chloramine  distils  in  the  form  of 
pale  yellow,  oily  drops  having  the  odor  of  nitrogen  chloride.  It  decom- 
poses slowly  in  dilute  solutions,  faster  in  concentrated  solutions,  yielding 
nitrogen,  ammonia,  and  hydrochloric  acid : 

3NH2C1=N2+NH3 


178  INORGANIC   CHEMISTRY.  [§§114- 

If  a  large  excess  of  ammonia  is  present,  it  acts  upon  the  chloramine  with 
the  formation  of  hydrazine  hydrochloride : 

NH2C1+NH3==NH2-NH2-HC1. 

The  addition  of  certain  substances,  e.g.,  a  small  amount  of  albumen, 
increases  the  yield,  giving  as  much  as  80%  of  the  theoretical  yield. 

By  fractional  distillation  of  the  aqueous  solution  the  hydrate 
N2H4-H20  is  obtained,  which  boils  constant  at  118.5°;  it  is  a  liquid  at 
ordinary  temperatures  and  freezes  below  —40°. 

LOBBY  DE  BRUYN  showed  that  the  molecule  of  water  can  be  removed 
by  treatment  with  barium  oxide  and  that  the  free  hydrazine  can  be 
obtained  in  the  pure  state  by  distillation  under  reduced  pressure.  This 
substance  is  liquid  at  ordinary  temperatures,  congeals  at  1.4°  and  boils* 
under  ordinary  pressure  at  113.5°.  Sp.g.  =  1.014  at  15°.  It  unites  with 
water  to  form  the  above  hydrate  with  the  evolution  of  heat.  Both  the 
free  hydrazine  and  its  aqueous  solution  have  a  strong  reducing  action. 
The  former  gradually  oxidizes  in  the  air,  reacts  vigorously  with  the 
halogens,  etc.  The  aqueous  solution  precipitates  the  metals  from 
solutions  of  salts  of  copper,  mercury,  silver,  etc.,  at  ordinary  tem- 
peratures. 

Hydrazine,  like  ammonia,  unites  with  acids  directly  to  form  salts;  it 
can  take  up  either  one  or  two  molecules  of  a  monobasic  acid,  N2H4  •  HC1 
and  N2H4-2HC1  being  both  known.  The  aqueous  solution  of  hydra- 
zine is  strongly  basic.  Its  salts  are  easily  soluble  in  water,  excepting 
the  sulphate,  N2H4-H2SO4,  which  is  rather  difficultly  so. 


HYDRAZOIC   ACID,  N3H. 

115.  This  interesting  compound,  like  the  preceding  one,  was  first 
discovered  by  CURTIUS  in  the  decomposition  of  an  organic  compound. 
It  can  now  be  prepared  in  a  good  yield  by  treating  hydrazine  hydrate 
with  nitrous  acid.  This  is  best  done  by  boiling  an  alcoholic  solution  of 
hydrazine  hydrate  with  amyl  nitrite  and  sodium  alcoholate,  which  gives 
the  sodium  salt  of  the  hydrazoic  acid. 

An  aqueous  solution  of  the  free  acid  is  best  obtained  by  distilling 
lead  hydrazoate,  Pb(N3)2,  with  dilute  sulphuric  acid.  By  fractional 
distillation  of  this  solution  the  pure  acid  can  be  obtained. 

Pure  hydrazoic  acid  is  a  liquid  with  a  penetrating,  unbearable  odor;  it 
boils  at  37°  and  is  extremely  explosive,  even  in  aqueous  solution. 

It  is  a  strange  fact  that  hydrazoic  acid  displays  more  or  less  analogy 
with  the  hydrogen  acids  of  the  halogens;  it  forms,  like  them,  difficultly 
soluble  salts  of  silver,  mercury  (ous)  and  lead.  These  are,  however, 


116.]    COMPOUNDS   OF  NITROGEN  WITH  THE  HALOGENS.   179 

soluble  in  strong  mineral  acids.  They  are  also  very  explosive,  hence 
extremely  dangerous,  the  sodium  salt  being  the  least  so.  An  aqueous 
1%  solution  of  the  acid  is  only  0.008  ionized;  it  is  thus  a  rather  weak 
acid;  it  gives  off  hydrogen  in  contact  with  many  metals,  e.g.  Zn,  Fe, 
Cd,  and  Mg.  It  is  characteristic  of  the  metal  hydrazoates  (or  "azides") 
that  they  crystallize  anhydrous  and  yield  the  pure  metal  when  heated. 

Compounds  of  Nitrogen  with  the  Halogens. 

116.  When  chlorine  gas  is  allowed  to  act  on  a  concentrated 
solution  of  ammonium  chloride,  most  conveniently  by  inverting  a 
flask  full  of  chlorine  over  the  warm  (30°-40°)  solution,  oily  drops 
are  formed,  which  are  best  collected  in  a  leaden  saucer  placed 
under  the  mouth  of  the  flask.  These  drops  contain  some  hydrogen 
as  well  as  nitrogen  and  chlorine.  By  treating  with  chlorine  once 
more  pure  nitrogen  trichloride,  NCls,  is  obtained  as  a  yellowish 
oil  with  a  disagreeable  pungent  odor  and  a  specific  gravity  of  1.65. 
This  is  one  of  the  most  dangerous  of  substances,  because  it 
explodes  in  a  most  violent  manner,  not  only  on  contact  with  certain 
organic  substances  (e.g.  turpentine),  but  very  often  spontaneously. 
It. dissolves  in  carbon  disulphide,  benzene  and  other  solvents,  form- 
ing yellow  solutions.  These  solutions  are  relatively  harmless;  they 
decompose  in  the  sunlight. 

Concentrated  hydrochloric  acid  decomposes  nitrogen  trichloride 
according  to  the  equation: 

NC13+4HC1=NH4C1+3C12; 
aqueous  .ammonia  also  breaks  it  up  in  a  similar  way: 

NC13 + 4NH3  =  3NH4C1 + N2. 
Nitrogen  trichloride  is  strongly  endo thermic : 
N+3C1-NC13  =  -41.9  Gal. 

When  a  solution  of  sodium  azide,  NaN3,  is  mixed  with  a  solution 
of  sodium  hypochlorite  in  the  relation  of  molecule  for  molecule  and  the 
mixture  is  acidified,  the  liquid  assumes  a  yellow  color  and  gives  off  a 
colorless  gas  with  an  odor  like  that  of  hypochlorous  acid  and  having 
the  composition  N3C1,  showing  it  to  be  chlorazide.  On  being  passed 


180  INORGANIC  CHEMISTRY.  [§§  116- 

into  caustic  soda  it  forms  sodium  azide  and  sodium  hypochlorite  in 
equivalent  amounts  : 

N,C1  +2NaOH  =  NaN3  +NaOCl  +H2O. 

The  chlorazide  is  likewise  extremely  explosive. 

117.  Nitrogen  Iodide.—  If  a  solution  of  iodine  in  potassium 
iodide  is  mixed  with  ammonia  solution,  a  precipitate  is  usually 
obtained  of  the  composition  NI2H;  if  the  conditions  are  slightly- 
altered  another  compound,  N2I3H3  (i.e.  NH3  +  NI3),  is  deposited 
which  breaks  up  on  continued  treatment  with  water  into  ammonia 
and  nitrogen  tri-iodide.  These  compounds  are  likewise  very 
explosive. 

Another  method  is  to  digest  pulverized  iodine  with  ammonia 
water.  The  product  so  obtained  is  still  more  explosive,  often 
exploding  even  when  damp  or  when  it  is  being  washed  with  water 
or  by  the  action  of  hydrochloric  acid.  In  the  presence  of  ammonia 
solution  it  is  stable. 

Nitrogen  iodide  is  decomposed  by  dilute  hydrochloric  acid, 
forming  ammonia  and  chlorine  iodide: 


Nitrogen  iodide  is  also  decidedly  endothermic. 


Hydroxylamine,  NH2OH. 

118.  Hydroxylamine  is  a  reduction  product  of  many  oxygen 
compounds  of  nitrogen  intermediate  to  the  formation  of  ammonia; 
e.g.  it  is  formed  when  tin  acts  on  dilute  nitric  acid.  Here  the 
nascent  hydrogen  effects  the  reduction: 

HN03  +  3H2  =  NH3O  +  2H20. 

It  is  manufactured  by  the  electrolytic  reduction  of  nitric  acid  dis- 
solved in  sulphuric  acid. 

The  free  hydroxylamine  is  best  prepared  by  heating  the  phosphate. 
It  is  a  crystallized  solid,  melting  at  30°  and  boiling  under  60  mm. 
pressure  at  70°.  When  heated  in  the  air  it  explodes  with  a  yellow 
flame. 


119.1  NITROUS  OXIDE.  181 

Hydroxylamine  is  easily  soluble  in  water;  its  solution  reacts 
strongly  alkaline.  It  forms  salts  in  the  same  way  as  ammonia,  i.e. 
by  direct  addition  of  the  acid:  NH2OH-HC1,  NH2OH-HN03,  etc. 
These  salts  are  rather  stable;  the  hydrochloride,  however,  must 
be  preserved  over  lime,  else  it  slowly  decomposes,  for  the  following 
reason.  The  salt  is  split  up  to  a  very  small  degree  into  hydrochloric 
acid  and  hydroxylamine.  Now  free  hydrochloric  acid  accelerates 
catalytically  the  decomposition  of  the  salt.  When,  however,  the 
hydrochloric  acid  is  absorbed  by  the  lime,  the  decomposition 
becomes  so  slow  that  it  is  imperceptible.  The  free  hydroxylamine 
and  its  aqueous  solution  are  somewhat  unstable,  especially  in  the 
presence  of  alkalies;  it  decomposes  easily  into  ammonia,  water  and 
nitrogen. 

A  further  characteristic  of  hydroxylamine  is  its  great  reducing 
power;  it  precipitates  reddish-yellow  cuprous  oxide  from  an  alkaline 
copper  solution  at  ordinary  temperatures,  even  when  strongly 
diluted;  mercuric  chloride,  HgCl2,  is  reduced  to  calomel,  Hg2Cl2; 
silver  nitrate  to  silver,  etc. 

The  following  reaction  is  also  peculiar :  A  solution  of  ferrous  sulphate 
is  precipitated  with  an  excess  of  sodium  hydroxide  and  warmed;  if 
hydroxylamine  (or  one  of  its  salts)  is  now  added  to  the  green  ferrous 
hydroxide,  red  ferric  hydroxide  is  formed  very  quickly,  the  hydrox- 
ylamine being  reduced  in  this  alkaline  solution  to  ammonia.  On  acidify- 
ing, an  acid  solution  of  a  ferric  salt  is  obtained;  if  this  is  treated  with 
a  hydroxylamine  salt,  it  is  suddenly  decolorized  because  of  reduction 
to  ferrous  salt,  the  hydroxylamine  being  now  in  the  oxidized  condition 
in  the  acid  solution. 

Compounds  of  Nitrogen  with  Oxygen. 

Those  included  under  this  title  are:  nitrous  oxide,  N2O;  nitric 
oxide,  NO;  nitrogen  trioxide,  or  nitrous  anhydride,  N203;  nitrogen 
dioxide,  NO2,  or  tetroxide,  N204,  and  nitrogen  pentoxide,  or  nitric 
anhydride,  N2Os. 

NITROUS  OXIDE,  N2O. 

119.  This  compound  cannot  be  obtained  directly  from  its  ele- 
ments; the  ordinary  method  of  preparation  consists  in  heating 
ammonium  nitrate  to  about  250°: 


182  INORGANIC  CHEMISTRY.  [§§  119- 

This  method  is  analogous  to  that  of  preparing  nitrogen  from 
ammonium  nitrite  (§  105).  If  the  nitrate  is  heated  above  250°, 
the  gaseous  product  partially  decomposes. 

Physical  Properties. — Nitrous  oxide  is  a  colorless  and  odorless 
gas;  which  when  liquefied  boils  at  —87°  and  solidifies  at  —102°. 
The  evaporation  of  the  liquid  produces  a  great  depression  in  the 
temperature,  which  may  even  reach  — 140°  under  reduced  pressure. 
Its  specific  gravity  is  1.52  (based  on  air),  or  21.89  for  0=16.  1  1. 
N20  at  0°  and  760  mm.  pressure  weighs  1.9657  g.  It  is  rather 
soluble  in  water  (1  vol.  H20  dissolves  1.305  vol.  N2O  at  0°);  hence 
it  must  be  collected  over  hot  water.  In  alcohol  it  is  still  more 
soluble. 

Chemical  Properties.  —  Nitrous  oxide  supports  combustion. 
Phosphorus,  carbon  and  a  glowing  splinter  burn  in  it  as  in  oxygen. 
A  mixture  of  nitrous  oxide  and  hydrogen  explodes  like  detonating- 
gas  when  it  is  ignited,  only  not  quite  so  loud.  These  properties 
might  lead  one  to  confuse  it  with  oxygen  on  a  superficial  examina- 
tion. However,  it  is  very  easily  distinguished  from  the  latter  by 
the  fact  that  it  gives  no  red  fumes  when  mixed  with  nitric  oxide 
(§  120)  and  always  leaves  residual  gas  (nitrogen)  after  a  combustion. 
A  faintly  burning  piece  of  sulphur  is  moreover  extinguished  by 
nitrous  oxide. 

Nitrous  oxide  is  endothermic:   2N+O  — N20=  — 17.7  Cal. 

BERTHELOT  has  made  the  general  observation  that  endothermic 
substances  can  suffer  an  explosive  decomposition;  in  this  case  this 
may  be  brought  about  by  touching  off  the  gas  with  fulminating 
mercury.  It  is  easy  to  explain  BERTHELOT'S  observation.  When 
an  endothermic  substance  decomposes,  heat  is  evolved.  Now,  we 
saw  in  §§  13  and  104  that  chemical  reactions  are  accelerated  in  a 
very  high  degree  by  rise  of  temperature.  Suppose  that  a  sudden 
decomposition  is  caused  at  a  certain  point  in  a  mass  of  an  endo- 
thermic compound.  The  heat  given  off  raises  the  temperature 
of  the  surrounding  molecules  and  they  too  split  up  suddenly, 
evolving  still  more  heat,  and  so  on.  The  whole  mass  will  thus 
reach  a  condition  of  sudden  decomposition,  that  is,  it  will  explode. 
To  bring  this  about  it  is  only  necessary  that  the  first  impulse  be 
vigorous  enough  for  the  sudden  decomposition  of  so  many  mole- 
cules that  the  heat  evolved  is  sufficient  to  raise  the  surrounding 
ones  to  the  temperature  of  decomposition. 


120.]  NITRIC  OXIDE.  183 

Composition.  —  Under  the  protracted  action  of  induction  sparks 
the  gas  splits  up  into  a  mixture  of  nitrogen  and  oxygen,  the  volume 
of  which  is  half  again  as  great  as  that  of  the  nitrous  oxide.  When 
potassium  and  sodium  are  burned  in  the  gas,  potassium  and  sodium 
oxides  respectively  are  formed,  together  with  nitrogen;  the  gas 
volume  after  cooling  is  unchanged.  Both  of  these  observations  point 
to  the  same  formula,  N2O,  and  this  is  confirmed  by  the  fact  that 
the  relative  density  of  the  gas,  which  should  theoretically  be 


was  found  to  be  21.89. 

NITRIC  OXIDE,  NO. 

1  20.  This  gas  is  only  obtained  by  the  reduction  of  nitric  or 
nitrous  acid.  The  ordinary  method  of  preparation  is  by  allowing 
copper  to  act  on  nitric  acid  or  else  by  covering  copper  (in  the  form 
of  thin  sheets)  with  a  saturated  solution  of  saltpetre  and  adding 
concentrated  sulphuric  acid  drop  by  drop  (§  127): 

3Cu  +  8HN03  =  3Cu(NO3)  2  +  4H20  +  2ND. 

In  this  reaction  the  hydrogen,  which  would  be  expected  to  be 
given  off,  reduces  another  portion  of  the  acid. 

In  order  to  prepare  nitric  oxide  by  the  reduction  of  nitric  acid 
or  a  nitrate  a  boiling-hot  solution  of  ferrous  chloride,  EeCl2,  in 
hydrochloric  acid  is  found  very  satisfactory;  the  ferrous  chloride 
is  converted  into  the  ferric  chloride,  FeCl3,  by  the  reaction: 

HN03  +  3FeCl2  +  3HC1  =  3FeCl3  +  2H20  +  NO. 

Perfectly  pure  nitric  oxide  is  obtained  by  treating  a  mixture  of  yellow 
prussiate  of  potash  and  potassium  nitrite  with  acetic  acid: 

2K4Fe(CX)6  +  2KNOa  +4C2H4O2  =K6Fe2(CN)12  +4KC2H3O2  +2H2O  +2NCX 

Yellow  prus-        Pot.  ni-      Acet.  acid.         Red  prus-  Pot.  acetate, 


siate.  trite 


Physical  Properties. — Nitric  oxide  is  a  colorless  gas,  whose 
specific  gravity  has  been  found  to  be  1.039  (air=l).  It  can  be 
condensed  to  a  blue  liquid,  which  boils  under  ordinary  pressure  at 


184  INORGANIC  CHEMISTRY.  [§§  120- 

—  153.6°.  The  critical  temperature  is  —  93.5°;  the  critical  pressure 
71.2  atm.  It  is  not  very  soluble  in  water,  but  dissolves  easily  in  a 
solution  of  ferrous  sulphate,  FeSCU;  strange  to  say,  this  solution  is 
quite  dark  brown  in  color,  although  the  ferrous  salt  solution  is  pale 
green  and  nitric  oxide  colorless.  The  compound  which  is  formed 
here  has  not  been  isolated,  but  it  has  been  shown  to  consist  of 
FeS04  and  NO  in  equimolecular  proportions. 

Chemical  Properties. — It  is  characteristic  of  this  gas,  above  all 
other  properties,  that  it  combines  with  oxygen  immediately, 
forming  nitrogen  dioxide,  a  reddish-brown  gas.  On  heating  it 
with  hydrogen  no  explosion  occurs;  the  mixture  burns  with  a 
white  flame,  forming  water  and  nitrogen.  If  burning  phosphorus 
is  introduced  into  the  gas,  it  continues  to  burn;  a  lighted  candle 
is,  however,  extinguished ;  sulphur  and  charcoal  do  not  burn  in 
it  either.  A  mixture  of  nitric  oxide  and  carbon  disulphide  burns 
with  an  intensely  luminous  blue  flame,  that  is  very  rich  in 
chemically  effective  rays. 

Nitric  oxide  is  a  strongly  endothermic  compound;  it  can  be 
made  to  explode  by  fulminating  mercury  (§  119).  According  to 
§  103  NO  must  be  formed  at  a  high  temperature.  NERNST 
proved  that  the  reaction  N2  +  02<=^2NO  accords  strictly  with  the 
law  of  mass-action  (§  49);  from  whichever  side  one  starts,  the 
results  are  in  agreement  with  those  calculated,  assuming  both 
reactions  to  be  bimolecular.  See  further  §  127. 

The  formation  and  the  decomposition  of  NO  are  much  slower 
than  in  the  case  of  ozone.  Accordingly  a  short  heating  of  air, 
followed  by  a  rapid  cooling,  produces  ozone,  while  a  slower 
heating  and  cooling  yield  NO,  ozone  being  broken  up  during  the 
extended  period  of  cooling.  The  following  experiment  illustrates 
this:  When  moist  air  is  directed  with  a  velocity  of  less  than  7  m. 
per  sec.  against  an  incandescent  NERNST  filament  (§  291)  NO  is 
formed;  a  more  rapid  current  gives  ozone. 

Composition. — When  sodium  is  heated  in  contact  with  a  measured 
amount  of  nitric  oxide,  sodium  oxide  and  nitrogen  are  formed ;  the 
latter  takes  up  exactly  half  the  volume  of  the  original  gas.  The 
specific  gravity  of  nitric  oxide  is  15  (H  =  l),  hence  its  molecular 
weight  is  30.  According  to  the  above  decomposition  the  gas  con- 
tains one  atom  of  nitrogen  (14  parts  by  weight).  There  remain 
for  the  oxygen,  therefore,  16  parts  by  weight,  i.e.  just  one  atom. 
Hence  the  formula  is  NO. 


122.]  NITROGEN  DIOXIDE  AND   TETROXIDE.  185 

Since  nitrogen  is  trivalent  or  quinquivalent  (the  latter  in  ammonium 
salts,  e.g.  NH4C1)  and  oxygen  is  bivalent,  it  must  be  assumed  that  there 
is  a  free  valence  bond  in  NO,  i.e.  -N=O.  The  same  applies  to  N02r 
Free  bonds  like  these  are  very  rare. 

Nitrous  Anhydride,  N20g. 

121.  Upon  allowing  nitric  acid  of  1.3  specific  gravity  to  react  with 
arsenic  trioxide  and  drying  very  carefully  the  gas  that  comes  off  and 
finally  condensing  this  gas,  a  liquid  is  obtained  of  the  composition  N203. 
At  ordinary  temperatures  the  liquid  is  green  but  below  —2°  it  is  deep 
indigo-blue.     When  cooled  by  liquid  air  it  solidifies  in  dark  blue  crystals  . 

2  x  14.  +3  v  16 
The  compound  N203  should  have  the  vapor  density  -     —  ^—t  -  =  38. 

In  a  perfectly  dry  state  its  vapor  density  was  found  to  vary  in  a  series 
of  experiments  between  38.1  and  62.2,  so  that  the  compound  appears 
in  that  condition  to  be  partly  polymerized.  The  least  trace  of  moisture, 
however,  causes  a  dissociation  into  NO2  and  NO.  For  instance,  it  was 
found  to  be  sufficient  merely  to  leave  a  small  bulb  full  of  it  with  a 
capillary  tube  open  a  few  seconds  in  the  air;  upon  resealing  the  tube 
the  vapor  density  was  found  to  have  fallen  to  28.2. 

This  very  striking  property  of  water,  whereby  even  the  slightest 
trace  of  it  brings  about  dissociations  which  are  not  observed  in  the 
perfectly  dry  state,  will  be  met  with  in  several  examples  in  later  chapters. 
The  phenomenon  was  discovered  by  BRERETON  BAKER. 

NITROGEN  DIOXIDE  AND  TETROXIDE,  N02  AND  N^04. 

122.  Nitrogen  dioxide  is  formed  from  nitric  oxide  plus  oxygen, 
or  more  convenient!^  by  heating  well-dried  lead  nitrate: 


When  so  prepared  it  is  a  very  deep-brown  gas.  On  leading  it 
into  a  strongly  cooled  vessel  it  condenses  to  a  bright-yellow  liquid, 
which  solidifies  at  —20°  to  colorless  crystals,  that  melt  at  —12°. 
The  color  becomes  darker  on  warming  and  at  +26°  the  liquid 
begins  to  boil,  changing  back  again  into  the  brown  gas.  The 
vapor  density  of  this  gas  at  26°  is  found  to  be  38.0,  while  that  cal- 
culated for  N2O4  is  45.9  and  that  for  NO2  22.9  (H  =  l).  ^  Since 
the  value  found  is  between  the  two,  "it  may  be  assumed  that  at  this 
temperature  the  vapor  consists  partly  of  N2O4  molecules  and  partly 
of  NO2  molecules.  A  simple  calculation  indicates  the  percentage 


186  INORGANIC  CHEMISTRY.  [§§  122- 

of  the  former  to  be  34.4%.  As  the  temperature  rises,  the  vapor 
density  steadily  decreases  till  about  150°  is  reached,  when  it  becomes 
constant  at  22.9.  There  is  evidently  complete  dissociation  of 
molecules  in  this  case, 


1  vol.          2  vols. 

and,  inasmuch  as  the  color  of  the  gas  grows  darker,  we  must  sup- 
pose that  NO2  is  dark  brown,  while  N2C>4  is  colorless,  which  is  true 
of  the  latter  in  the  solid  state.  This  supposition  is  supported  by 
the  fact  that  not  only  can  the  degree  of  dissociation  be  estimated 
from  the  intensity  of  the  color,  but  that  it  can  even  be  measured 
quantitatively  in  this  way. 

According  to  §  51  the  equilibrium  between  the  two  gases  is 
expressed  by  the  equation 

P-x=kx2, 

where  P  is  the  total  pressure  of  the  gas  mixture  and  x  that  of  the 
dioxide,  k  being  a  constant.  From  this  equation  it  follows  that 
the  dissociation  (at  a  constant  temperature)  depends  on  the  pres- 
sure (§51),  which  has  been  shown  to  be  the  case.  This  also  fol- 
lows from  the  theorem  of  LE  CHATELIER  (§  102). 

On  bringing  nitrogen  tetroxide  in  contact  with  water  or,  better, 
with  alkalies,  nitrous  and  nitric  acids  are  formed;  we  may  therefore 
consider  it  as  a  mixed  anhydride  of  these  two  acids  : 


Nitric  acid.      Nitrous  acid. 

Both  N02  and  N204  possess  strong  oxidizing  power;  many 
substances  burn  in  their  vapor;  they  precipitate  iodine  from  solu- 
ble iodides. 

The  composition  of  nitrogen  dioxide  follows  from  its  synthesis- 
equation,  2NCM-02,  and  from  the  vapor  density. 

Nitrogen  Pentoxide,  N205. 

123.  This  compound  can  be  obtained  by  the  action  of  chlorine  on 
silver  nitrate  or  by  distilling  fuming  nitric  acid  with  phosphorus  pent- 
oxide.  It  is  a  colorless  crystalline  solid.  It  melts  at  30°,  and  at  45-50° 
breaks  up,  giving  off  brown  fumes.  If  the  heating  takes  place  rather 


125.]  HYPONITROUS  ACID.  187 

rapidly  the  decomposition  is  explosive  in  nature;  sometimes  a  spon- 
taneous explosion  takes  place,  hence  it  can  not  be  kept  long. 

As  nitrogen  pentoxide  is  strongly  endothermic,  its  spontaneous 
explosion  must  be  explained  in  the  same  way  as  is  indicated  in  §  119. 
Only  we  must  conclude  in  this  case  that  the  decomposition  at  ordinary 
temperatures  is  vigorous  enough  to  sufficiently  heat  the  neighboring 
molecules. 

It  unites  with  water,  forming  nitric  acid  with  the  evolution  of  much 
heat.  As  might  be  expected,  it  has  strongly  oxidizing  properties.  Phos- 
phorus and  potassium,  for  instance,  burn  with  great  brilliance  in  the 
slightly  warmed  anhydride. 

The  composition  of  nitrogen  pentoxide  is  ascertained  by  heating  with 
powdered  copper;  the  amount  of  nitrogen  evolved  corresponds  to  the 
formula  N2O5 

Oxygen  Acids  of  Nitrogen. 

124.  Four    acids  of  nitrogen   are  known:    hyponitrous  acid, 
H2N2O2;  nitrohi/droxi(laminic  acid,  H2N203;  nitrous  acid,  HNO2; 
nitric  acid,   HNO3.     The  nitrous  acid  is  known  only  in  dilute 
aqueous  solution;   mtrohydroxylaminie  acid  is  known  only  in  its 
salts;  but  the  others  are  known  in  the  pure  state. 

Only  certain  ones  of  the  above  nitrogen  oxides  can  be  regarded 
as  acid  anhydrides.  The  pentoxide  is  undoubtedly  one  and  the 
tetroxide  may  be  considered  as  a  mixed  anhydride  of  nitric  and 
nitrous  acids  (§  122).  Nitrogen  trioxide  gives  a  solution  of  nitrous 
acid  when  mixed  with  water  at  a  low  temperature;  however,  this 
solution  undergoes  a  decomposition  slowly  at  ordinary,  more  rapidly 
at  higher,  temperatures;  nitric  acid  and  nitric  oxide  being  formed: 

3HN02=HN03+2NO  +  H2O. 

The  acid  corresponding  to  nitric  oxide,  NO,  is  nitrohydroxyl- 
aminic  acid.  However  no  one  has  yet  been  able  to  obtain  this 
acid  from  nitric  oxide  and  water.  The  same  is  true  for  nitrous 
oxide,  to  which  hyponitrous  acid  corresponds. 

Hyponitrous  Acid,  H2N202. 

125.  This  acid  is  formed  when  nitrogen  trioxide  is  introduced  into  a 
methyl-alcoholic  solution  of  hydroxylamine.     The  free  acid  does  not 
liberate  iodine  from  potassium  iodide  at   once;   the  reaction  is  delayed 


188  INORGANIC  CHEMISTRY.  [§§  125- 

for  a  time,  probably  on  account  of  a  decomposition,  by  which  nitrous 
acid  is  formed. 

Hyponitrous  acid  belongs  to  the  class  of  weak  acids;  its  aqueous 
solution  is  a  poor  conductor.  Both  neutral  and  acid  salts  of  this  acid 
are  known. 

Nitrohydroxylaminic  Acid,  H2N208. 

This  acid  does  not  exist  in  the  free  state,  being  known  only  in  salts. 
Its  sodium  salt  is  obtained  by  mixing  an  alcoholic  solution  containing 
sodium  alcoholate  and  hydroxylamine  with  ethyl  nitrate: 

C2H5ONO2  +  NH2OH  =  C2H5OH  +  H2N2O3. 

Ethyl  Nitrate 

The  alcoholate  is  added  in  order  to  convert  the  free  acid  directly  into  its 
sodium  salt.  If  the  attempt  is  made  to  liberate  it  by  adding  a  stronger 
acid,  it  is  immediately  decomposed  according  to  the  equation: 

Na«NA  +  2HC1  =  2NaCl  +  2NO  +  H20. 

The  sodium  salt,  heated  in  aqueous  solution,  gives  sodium  nitrate  -and 
nitrous  oxide.  When  the  sodium  salt  is  heated  dry  until  it  begins  to 
melt,  it  is  decomposed  into  nitrite  and  hyponitrite: 

2Na2N203  =  2NaNO2  +  Na2N202. 


NITROUS  ACID,  HN02. 

126.  It  was  remarked  above  that  this  acid  is  only  known 
in  dilute  solution  at  ordinary  or  low  temperatures;  its  salts 
are,  however,  stable.  In  order  to  prepare  them  we  usually 
employ  potassium  or  sodium  nitrate,  which  gives  off  oxygen 
when  heated  and  is  converted  into  nitrite.  This  decomposition 
takes  place  more  readily  if  lead  is  added  during  the  heating 
as  a  reducing  agent: 

2KN03  =  2KN02  +  02. 


127.J  NITRIC   ACID.  ]8g 

Its  salts  are  all  easily  soluble  in  water,  "with  the  exception  of 
silver  nitrite,  AgN02,  which  is  rather  difficultly  soluble  at  ordinary 
temperatures ;  it  is  obtained  as  a  yellow  crystalline  precipitate,  when 
not  too  dilute  solutions  of  silver  nitrate  are  mixed  with  a  nitrite. 

The  addition  of  strong  sulphuric  acid  to  a  nitrite  at  once  pro- 
duces red  fumes;  in  this  way  a  nitrite  can  be  distinguished  from 
a  nitrate,  for  the  latter  does  not  produce  them.  It  may  be  assumed 
that  in  this  reaction  free  nitrous  acid  is  primarily  formed;  this  is, 
however,  broken  up  directly  into  water  and  nitrogen  trioxide,  the 
latter  of  which  at  once  splits  up  again  into  NO2  +  NO;  thereupon 
the  nitric  oxide  unites  immediately  with  the  surrounding  oxygen 
to  form  dioxide.  The  red  fumes  thus  consist  solely  of  nitrogen 
dioxide,  N02. 

On  treating  a  very  dilute  nitrite  solution  with  the  equivalent 
amount  of  sulphuric  acid  a  dilute  solution  of  free  nitrous  acid  is 
obtained.  This  solution  can  act  either  oxidizing  or  reducing.  As 
examples  of  the  former  action  we  have  the  liberation  of  iodine  from 
a  solution  of  potassium  iodide,  the  oxidation  of  sulphurous  acid  in 
dilute  solution  to  sulphuric  acid,  the  oxidation  of  ferrous  sul- 
phate, FeS04,  to  ferric  sulphate,  Fe2(SO4)3,  and  the  conversion  of 
the.  yellow  to  the  red  prussiate  of  potash.  In  all  of  these  cases  lower 
oxides  of  nitrogen,  chiefly  nitric  oxide,  are  formed.  An  example 
of  its  reducing  action  (in  which  nitrous  acid  is  oxidized  to  nitric 
acid)  is  the  bleaching  of  potassium  permanganate,  KMn04,  in  sul- 
phuric acid  solution: 

2KMnO4  +  5HNO2  +  3H2S04  =  K2SO4 + 2MnS04  +  5HNO3  +  3H20. 

This  last  reaction  offers  a  means  of  determining  quantitatively 
(volumetrically,  see  §  93)  the  strength  of  a  dilute  solution  of  nitrous 
acid. 

NITRIC  ACID,  HN03. 

127.  This  is  the  best  known  acid  of  nitrogen.  It  is  manufac- 
tured on  a  large  scale,  since  its  uses  are  many  and  varied;  in  the 
organic  dyestuff  industry,  for  example,  large  quantities  are  employed. 
The  commercial  process  of  manufacture  depends  on  the  decom- 
position of  Chili  saltpetre,  NaN03,  by  strong  sulphuric  acid: 

NaN03  +  H2S04  =  NaHS04 + HN03 

One  of  the  simplest  methods  of  carrying  it  out  is  as  follows: 


190 


INORGANIC  CHEMISTRY. 


127- 


In  the  cast-iron  .retort  (C,  Fig.  35),  saltpetre  and  sulphuric  acid 
(chamber-acid)  are  mixed  in  proportions  corresponding  to  the  above 
equation,  a  slight  excess  of  sulphuric  acid,  however,  being  added, 
because  this  makes  the  residue  easier  to  remove  from  the  retort, 
The  retort  is  connected  with  a  row  of  earthenware  bottles  (EEf) 
containing  a  little  water.  These  receive  the  distilled  acid.  The 
last  bottle  connects  with  a  coke  tower  through  which  water  is 


FIG.  35.— MANUFACTURE  OF  NITRIC  ACID. 

trickling  down  to  dissolve  the  uncondensed  acid  vapor.  By  this 
process  a  liquid  of  a  specific  gravity  of  1.35  and  containing  60% 
acid  is  obtained.  If  the  saltpetre  is  previously  dried  and  concen- 
trated sulphuric  acid  is  used,  a  nitric  acid  of  sp.  g.  1.52  and  almost 
100%  pure  can  be  obtained. 

In  some  cases  two  molecules  of  saltpetre  are  used  to  one  of 
sulphuric  acid.  If  heat  is  moderately  applied,  the  reaction  pro- 
ceeds according  to  the  above  equation,  but  on  heating  to  a  higher 
temperature  the  acid  sodium  sulphate  that  is  formed  acts  on  the 
second  molecule  of  nitrate,  also  forming  nitric  acid: 

NaN03  4-  NaHS04  =  Na2S04  +  HN03. 

A  large  part  of  the  nitric  acid,  however,  dissociates  at  the  same 
time  as  follows: 

2HN03=2N02-fH20  +  0. 

The  N02-fumes  dissolve  in  the  distillate.  The  liquid  thus 
obtained  is  red  and  its  specific  gravity  is  1.52-1.54;  it  fumes 
strongly  in  the  air  and  is  known  as  "red  fuming  nitric  acid." 


127.] 


NITRIC  ACID. 


191 


For  some  years  the  distillation  of  saltpetre  with  sulphuric  acid 
has  been  carried  on  ;n.  a  vacuum.  The  yield  of  acid  in  such  a  case 
approaches  closely  to  the  theoretical  and  the  product  obtained  is 
entirely  free  from  nitrous  fumes. 

An  entirely  distinct  method  for  the  industrial  preparation  of 
nitric  acid  was  Lay:nted  a  few  years  ago  by  BIRKELAND  and 
EYDE.  They  make  use  of  the  nitrogen  and  oxygen  of  the  atmos- 
phere. The  problem  of  making  nitric  acid  from  this  rather  in- 
exhaustible source  has  been  studied  for  many  years,  but  these 
men  are  the  first  to  handle  it  with  success  on  a  commercial  scale. 
The  solution  of  the  problem  became  a  really  pressing  matter, 
because  the  principal  material  Tor  the  preparation  of  nitrogen 
compounds,  Chili  saltpetre,  bids  fair  to  be  exhausted  in  thirty- 
five  years,  and  saltpetre  has  not  only  a  large  significance  in  the 
industrial  world  but  a  still  larger  one  in  agriculture  as  a  nitro- 
genous fertilizer.^ 

The  method  of  BIRKELAND  and  EYDE  is  based  on  the  long 
established  fact  that  oxides  of  nitrogen  are  formed  in  an  electric 
arc  burning  in  the  air.  The 
reason  why  previous  investiga- 
tions did  not  succeed  lies  in 
the  fact  that  an  ordinary  electric 
arc  has  too  small  a  volume,  and 
therefore,  cannot  let  a  sufficient 
quantity  of  air  pass.  This  dif- 
culty  is  now  obviated  by  mount- 
ing the  arc  between  the  poles 
PP  of  a  very  powerful  electric 
magnet  EE  (Fig.  36).  The  arc 
is  produced  between  two  hol- 
low bars  of  copper,  which  are 
kept  cool  by  circulating  water 
in  them.  When  an  alternate 
current  is  used  for  producing 
the  arc,  the  latter  spreads  out  FIG.  36.— DIAGRAM  OF  BIRKELAND  AND 
in  the  shape  of  a  flat  disc  that  EYDE  NITRIC  ACID  APPARATUS. 
reaches  a  diameter  of  2  m.  in 

the  industrial   form   of  the   apparatus;    the   tension  employed 
is  5000  volts.     This    flame  disc  is  inclosed  in  a  box  through 


192  INORGANIC  CHEMISTRY.  [§  127. 

which  a  rapid  current  of  air  is  forced,  and  the  contact  with 
the  flame  is  sufficient  to  form  somewhat  more  than  1%  of  NO. 
Instead  of  broadening  out  the  electric  arc  to  a  sun-shaped  disc 
by  the  action  of  powerful  magnets,  SCHONHERR  (Badische  Anilin- 
und  Sodafabrik)  forms  an  arc  in  the  inside  of  an  iron  pipe 
through  which  air  is  passed.  Under  these  circumstances  the 
arc  is  developed  in  a  peculiar  manner.  When  the  current  is 
turned  on,  the  arc  forms  at  the  first  instant  in  the  lower  part 
of  the  metal  pipe,  between  the  pipe  itself,  which  serves  as  an 
electrode,  and  a  second  electrode,  which  is  separated  by  only 
a  few  millimeters  from  the  lower  end  of  the  pipe.  Forthwith, 
however,  the  arc  is  carried  along  upward  in  the  pipe  by  the 
current  of  air,  which  is  given  a  tangential  motion  as  it  is  passed 
into  the  pipe,  so  that  the  arc  comes  to  occupy  the  portion  of 
air  along  the  axis  of  the  pipe  and  does  not  touch  the  wall  of 
the  pipe  (or  the  efflux  end  of  the  pipe  or  a  specially  devised 
separate  electrode)  until  a  considerable  distance  from  the  lower 
electrode  is  reached.  Thus  there  is  established  in  the  axis  of 
the  pipe  a  continuous  and  quietly  burning  column  of  light  of 
very  powerful  actinic  effect.  In  this  long-drawn-out  arc  the 
passing  air  is  partially  transformed  into  nitric  oxide.  This  is 
quickly  chilled  by  contact  with  the  wall  of  the  pipe,  which  is- 
externally  exposed  to  the  atmosphere,  and  so  prevented  from 
redecomposition.  The  gaseous  product  is  half  again,  if  not 
twice,  as  rich  in  nitric  oxide  (yield  about  2%)  as  by  the  BIRKE- 
LAND-EYDE  process. 

The  NO  must  be  looked  upon  as  the  primary  product,  which 
subsequently  unites  with  oxygen  to  form  NO2,  the  latter  being 
carried  to  water-absorption  towers  much  like  the  GAY-LUSSAC 
towers  in  sulphuric  acid  plants.  The  N02  cannot  be  the  primary 
product,  for  it  dissociates  at  about  600°  into  NO  and  02.  With 
the  water  the  NO2  forms  nitric  acid  and  nitrous  acid : 

N204  +  H20  =  HNO3  +  HN02. 
The  latter  yields  NO2  and  NO,  however,  when  the  liquid  becomes 


more  concentrated: 


2HNO2  =  H20  +  NO2  +  NO. 


§  127.]  NITRIC  ACID.  193 

NO  is  once  more  converted  into  NC>2  and  the  N02  again  gives 
nitric  acid;  eventually  all  is  converted  into  that  acid.  Instead 
of  marketing  sodium  nitrate,  to  duplicate  the  Chili  saltpetre, 
calcium  nitrate  is  produced  by  saturating  the  nitric  acid  with 
lime  and  the  resulting  calcium  nitrate  is  used  for  fertilizing  and 
other  purposes. 

Nitrites  are  also  manufactured  directly  by  leading  N2O4  into 
caustic : 

N2O4  +  2KOH  =  KN03  +  KN02  +  H2O. 

The  nitrate  and  nitrite  are  separable  by  fractional  crystallization. 

Physical  Properties. — Absolute  nitric  acid,  i.e.  the  compound 
HNO3  in  the  pure  state,  is  prepared  by  distilling  the  nearly  pure 
acid  of  commerce  (sp.  g.  1.5)  with  concentrated  sulphuric  acid 
in  vacuo.  The  liquid  distillate  has  a  specific  gravity  of  1.559  at 
0°  and  becomes  solid  at  —40°;  it  boils  under  ordinary  pressure 
at  86°,  but  with  partial  decomposition. 

Chemical  Properties. — Nitric  acid,  especially  when  pure,  is  a 
rather  unstable  compound;  at  ordinary  temperatures  it  is  decom- 
posed by  sunlight  to  a  slight  extent,  turning  yellow  on  account 
of  the  small  amount  of  nitrogen  dioxide  formed.  At  an  elevated 
temperature  the  acid  also  breaks  up,  decomposition  into  nitrogen 
dioxide,  water,  and  oxygen  being  complete  at  260°. 

When  strong  nitric  acid  is  subjected  to  repeated  distillation  under 
atmospheric  pressure,  its  boiling-point  gradually  rises,  while  the  acid 
becomes  proportionately  weaker,  until  finally  a  68%  acid  is  obtained, 
which  boils  constant  at  120.5°.  The  same  mixture  is  obtained  when 
one  starts  with  dilute  acid  and  distils  it.  In  both  cases  the  boiling- 
point  of  the  original  liquid  is  lower  than  that  of  the  product;  it  rises 
during  the  boiling  to  a  maximum  at  120.5°.  We  have  here,  therefore 
the  case  of  a  liquid  mixture  with  a  maximum  boiling-point,  which  is 
discussed  in  ORG.  CHEM.,  §  22.  The  mixture  of  hydrogen  chloride  and 
water  also  has  a  maximum  boiling-point  (110°). 

Nitric  acid  is  very  extensively  ionized  in  aqueous  solution; 
it  is  one  of  the  strongest  acids  known. 

When  it  comes  in  contact  with  metals,  the  salts  of  nitric  acid 
(nitrates)  are  formed,  but  without  any  evolution  of  hydrogen, 
since  part  of  the  acid  present  is  reduced  by  the  nascent  hydrogen. 
The  nitrates  are  all  easily  soluble  in  water.  The  action  of  nitric 


194  INORGANIC  CHEMISTRY.  [§§  127- 

acid  on  the  metals  is  not  the  same  in  all  cases.  It  does  not  attack 
gold  or  platinum.  Silver,  mercury,  and  copper  are  only  imper- 
ceptibly dissolved  at  ordinary  temperatures,  but  on  warming  they 
dissolve  with  the  evolution  of  nitric  oxide.  This  and  the  other 
NO-compounds  are  powerful  catalyzers  in  the  dissolving  of  the 
above-named  metals,  for  nitric  acid  which  is  perfecty  free  from 
them  does  not  dissolve  these  metals,  while  the  reaction  immediately 
begins  as  soon  as  a  little  of  these  substances  is  added:  It  may  be 
supposed  that  on  warming  nitric  acid  traces  of  NO-compounds 
are  formed,  which  together  with  the  elevation  of  the  temperature 
accelerate  the  reaction.  Iron,  zinc,  and  magnesium  reduce  nitric 
acid  to  nitrous  oxide  and  even  to  ammonia.  Under  the  action  of 
iron  filings  and  dilute  sulphuric  acid  the  reduction  of  nitric  acid 
to  ammonia  in  dilute  solution  is  quantitative.  There  are  also 
various  denitrifying  bacteria  known,  Bacillus  pyocyaneus  being  the 
best  studied  of  them. 

Nitric  acid  frequently  acts  as  a  powerful  oxidizing  agent, 
especially  at  an  elevated  temperature  If  sulphur  is  boiled  with  it, 
the  sulphur  is  converted  to  sulphuric  acid,  similarly  phosphorus  to 
phosphoric  acid.  A  glowing  piece  of  charcoal  dropped  upon  the 
concentrated  acid  continues  to  burn  with  a  bright  glow.  In  all 
these  cases  the  highest  oxidation  stages  are  formed.  Nitric  acid 
is  used  particularly  in  the  organic  branches  of  chemical  industry. 

The  composition  of  nitric  acid  can  be  deduced  from  that  of  its 
anhydride.  A  weighed  amount  of  the  latter  is  introduced  into 
water;  nitric  acid  is  formed,  which  is  neutralized  with  baryta 
water.  By  evaporation  it  is  possible  to  determine  how  many 
parts  by  weight  of  barium  oxide,  BaO,  combine  with  the  anhydride. 
It  is  found  that  153.37  parts  (=!BaO)  combine  with  108.02 
parts  (=1N2O5)  of  the  anhydride;  the  formula  of  barium  nitrate 
thus  becomes  Ba(NO3)2,  hence  that  of  nitric  acid  itself  imjist  be 
HN03. 

Pernitric  Acid,  HN04. 

Pernitric  acid  is  formed  in  very  dilute  aqueous  solution  by  the 
oxidation  of  nitrous  acid  with  hydrogen  peroxide : 

2H2O2  +  HN02  =  HN04  +  2H2O. 

Nitric  acid  does  not  yield  it  when  treated  in  the  same  way;  on  the 
contrary,  the  pernitric  acid  breaks  up  even  in  a  cold  dilute  aqueous 


128.]  DERIVATIVES  OF  THE  NITROGEN  ACIDS.  195 

solution  inside  of  about  an  hour  completely  into  nitric  acid  and  hydrogen 
peroxide  : 

HNO4  +  H20  -  HNO3  +  H202. 

Pernitric  acid  has  the  very  characteristic  property  of  liberating  bromine 
from  potassium  bromide  solutions,  something  that  neither  hydrogen 
peroxide  nor  nitrous  acid  nor  nitric  acid  does. 

Derivatives  of  the  Nitrogen  Acids. 

128.  In  discussing  the  manufacture  of  sulphuric  acid  (§  86) 
we  already  referred  to  the  chamber  crystals,  HSO5N.  They  are 
formed  in  the  lead  chambers  in  case  not  enough  steam  is  supplied. 
The  following  equation  expresses  the  action  that  takes  place: 

2S02  +  N2O4  +  0  +  H2O  =  2SO5NH. 

The  ordinary  method  of  preparing  this  substance  is  by  conduct- 
ing carefully  dried  sulphurous  oxide  into  cooled  fuming  nitric  acid: 

S02  +  HNO3=SO5NH. 

The  crystalline  mass  obtained  is  spread  out  on  porous  earthenware 
to  allow  the  adhering  liquid  to  be  absorbed. 

The  chamber  crystals  have  the  appearance  of  a  coarse  crystal- 
line, colorless  mass;  they  melt  at  73°.  They  are  at  once  decom- 
posed by  water  into  sulphuric  and  nitrous  acids: 

SO5NH  +  H2O  =  H2SO4  +  HN02. 

For  this  reason  the  compound  is  considered  as  the  mixed  anhy- 
dride of  sulphuric  and  nitrous  acids.  According  to  §  90  the  struc- 

OTT 

ture  S02<Qg  can  be  ascribed  to  sulphuric  acid;  to  nitrous  acid  the 

structure  HO-NO,  since  a  hydroxyl  group  is  assumed  (§  129)  to 
exist  in  it.  For  the  chamber  crystals  we  therefore  have 

OH 


Chamber  crystals. 


Since  the  atomic  group  NO  is  known  as  nitrosyl,  the  rational 
name  for  the  compound  is  nitrosyl  sulphuric  acid. 


196  INORGANIC  CHEMISTRY.  [§§  128- 

Nitrosyl  sulphuric  acid  plays  a  very  important  role  in  the 
lead-chamber  process.  In  its  decomposition  by  water  it  forms 
sulphuric  and  nitrous  acids,  according  to  the  equation  given 
above.  Nitrous  acid  reacts  with  sulphur  dioxide,  forming 
nitroso-sulphonic  acid : 

HN02  +  S02=ON-SO3H, 

which  is  very  unstable,  however,  and  unites  at  once  with  a 
second  molecule  of  nitrous  acid  to  form  nitrosi-sulphonic  acid 
and  nitric  oxide: 

/OH 

ON  -  S03H + HNO2 = ON<  +  NO ; 

XS03H 

but  the  nitrosi-sulphonic  acid,  likewise  very  unstable,  soon  splits 
up  under  the  conditions  prevailing  in  the  lead  chambers  into 
nitric  oxide  and  sulphuric  acid: 

/OH 

ON<  =NO  +  H2S04. 

\S03H 

The  nitric  oxide,  finally,  oxidizes  very  quickly  in  the  oxygen 
always  present  in  the  lead  chambers,  back  again  into  nitrous 
acid: 

2NO  +  O  +  H20  =  2HNO2. 

These  four  equations  explain  the  whole  lead-chamber  process, 
according  to  RASCHIG.  He  succeeded  in  preparing  the  nitroso- 
and  nitrosi-sulphonic  acids;  see  §  130. 

Concentrated  sulphuric  acid  dissolves  nitrosyl  sulphuric  acid 
without  change.  This  solution  is  very  stable;  it  can  be  distilled 
without  decomposition.  It  is  formed  in  the  Gay-Lussac  tower  of 
the  sulphuric  acid  factory  and  is  called  "  nitrated  acid"  or 
"  nitrous  vitriol."  On  dilution  with  water  this  solution  is  not 
altered  until  its  specific  gravity  reaches  1.55-1.50  (51°-48°B.); 
then  the  nitric  oxides  begin  to  escape,  especially  when  warmed. 
Water  and  sulphurous  oxide  act  on  the  nitrated  acid,  producing 
the  following  reaction: 


2SO2  <  Q^ + 2H2O  +  S02  =  3SO2  <  QH  + 2NO- 
This  action  takes  place  in  the  Glover  tower. 


129.]  DERIVATIVES  OF  THE  NITROGEN  ACIDS.  197 

It  is  of  great  importance  industrially  to  be  able  to  determine  the 
amount  of  nitrogen  that  the  nitrated  acid  contains.  This  can  be  done 
as  follows:  The  acid  is  agitated  with  mercury,  whereupon  all  nitrogen 
compounds  in  solution  are  given  off  in  the  form  of  nitric  oxide,  and 
the  gas  evolved  is  measured.  Another  method  consists  in  decomposing 
the  nitrated  acid  with  a  large  excess  of  water  and  titrating  the  resulting' 
nitrous  acid  with  permanganate  (§  126). 

129.  Nitrosyl  chloride,  NOC1,  is  a  reddish-yellow  gas  at  ordi- 
nary temperatures;  by  cooling  it  can  be  condensed  to  a  red  liquid, 
that  boils  at  —5.6°.  Its  melting  point  lies  at  —60°.  Its  critical 
temperature  is  164°.  This  compound  is  formed  by  the  direct  com- 
bination of  nitric  oxide  and  chlorine.  Bone-black  is  a  good  cata- 
lyzer of  this  reaction  ,  making  this  method  of  preparation  a  very 
satisfactory  one.  Another  method  is  that  of  carefully  warming 
chamber  crystals  with  thoroughly  dried  sodium  chloride  : 


SO  2  <          +  NaCl  =  S02  <          +  NOC1. 


Nitrosyl  chloride  is  broken  up  by  water  into  nitrous  and  hydro- 
chloric acids: 

NOC1  +  H20  =  NO.OH+HC1; 

hence  it  may  be  regarded  as  the  chloride  of  nitrous  acid  (§  130). 

Only  when  the  temperature  exceeds  700°  is  there  any  dissocia- 
tion into  NO  and  C12.  Aluminium  is  not  attacked  by  it  until  500° 
is  reached  ;  however,  silver  powder  decomposes  it  quantitatively  at 
this  temperature  into  nitric  oxide  and  chlorine,  the  latter  forming 
silver  chloride  with  silver.  This  method  is  used  for  the  analysis  of 
the  gas. 

On  mixing  hydrochloric  acid  with  nitric  acid  a  liquid  is  pro- 
duced which,  in  addition  to  these  two  acids,  contains  chlorine  and 
nitrosyl  chloride: 

HN03  +  3HC1  =  2H2O  +  NOC1  +  C12. 

This  liquid   (in  virtue  of  the  free  chlorine  it  contains)   dissolves 
the  precious  metals,  including  gold,  the  "king"  of  the  metals,  and 


198    '  INORGANIC  CHEMISTRY.  [§§  129- 

therefore  bears  the  name  aqua  regia.  It  was  known  even  to  the 
alchemists,  particularly  GEBER,  who  prepared  it  by  dissolving  sal 
ammoniac,  NE^Cl,  in  nitric  acid. 


OTHER  NITROGEN  COMPOUNDS. 

130.  In  organic  chemistry  acid  derivatives  are  known  which  are 
formed  through  the  substitution  of  the  hydroxyl  group  by  a  halogen 
atom  or  the  amido  group,  NH2.  The  former  class  are  called  acid 
chlorides,  the  latter  amides.  Of  the  inorganic  acid  chlorides 
we  have  already  become  acquainted  with  those  of  sulphuric  acid  (§90). 
Nitrosyl  chloride  is  also  an  acid  chloride. 

Upon  mixing  concentrated  solutions  of  sodium  nitrite  and  sodium 
bisulphite  the  liquid  heats  strongly  and  a  reaction  takes  place  according 
to  the  following  equation: 

NaN02  +  SNaHSO,  =  N  (S03Na) , + NaOH  +  H20. 

The  resulting  compound,  to  which  has  been  assigned  the  structural 
formula  N(S03Na)3  +  2H2O  after  RASCHIG,  is  regarded  as  ammonia  in 
which  three  hydrogen  atoms  have  been  replaced  by  sulpho-groups 
(S03H) ;  it  bears  the  name  sodium  nitrilo-sulphonate.  It  is  very  soluble 
in  water;  but,  if  the  liquid  mixture  is  poured  into  a  cold-saturated 
potassium  chloride  solution,  the  difficultly  soluble  N(S03K)3  is  pre- 
cipitated. On  boiling  it  for  a  short  time  with  water,  or  better,  by  letting 
it  stand  for  a  day  moistened  with  dilute  sulphuric  acid,  it  forms 
potassium  imidosulphonate: 

N  (S03K)  3  +  H20  =  NH  (S03K)  2 + KHS04. 

If  the  boiling  is  continued  for  a  long  period,  amidosulphonic  acid  is  pro- 
duced : 

N  (SO3K)3  +  2H20  =  NH2SO3H  +  K2S04  +  KHS04. 

Similar  derivatives  of  hydroxylamine  are  also  known.  The  potassium 
salt  of  hydroxylamine-disulphonic  acid,  HO-N(S03K)2+2H2O,  is  formed 
when  solutions  of  acid  potassium  sulphite  and  potassium  nitrite  are 
mixed : 

KNO2  +2KHSO3  =HON(SO3K)2  +  KOH. 

By  boiling  this  compound  with  water  hydroxylamine  is  set  free: 
HON(SO3K)2  +2H,O  =HONH2  +  2SO4KH. 


131.]  PHOSPHORUS.  199 

Oxidation  of  sodium  hydroxylamine-disulphonate  with  permanganate 
solution  gives  an  intensely  violet-colored  liquid.  This  coloration  is  due 
to  the  Na  salt  of  the  nitrosi-sulphonic  acid  mentioned  in  §  128.  Addi- 
tion of  cold-saturated  potassium  chloride  solution  causes  the  separation  of 
orange  crystals  of  the  corresponding  potassium  salt,  which  is  also  soluble 
in  water  with  a  violet  color. 

If  a  solution  of  chamber  crystals  in  concentrated  sulphuric  acid  is 
shaken  with  mercury,  the  solution  becomes  sky-blue  on  account  of 
the  reduction  of  the  nitrosyl  sulphuric  acid  to  nitrosi-sulphonic  acid, 


.     The  copper  salt  of  the  latter  is  more  stable  and  can  be 
SO3H 

obtained  by  shaking  the  above-mentioned  solution  with  copper  foil 
clippings.  The  addition  of  water  or  oxidizing  agents  (e.g.,  nitric  acid) 
destroys  the  blue  coloration. 

The  amide  of  nitric  acid,  nitramide,  NH2-N02,  has  so  far  been  obtained 
only  from  an  organic  compound,  nitrourethane,  N02NH  •  C02C2H5. 
This  amide  appears  in  colorless  crystals  that  melt  at  72°  with  decom- 
position. The  aqueous  solution  reacts  strongly  acid.  It  is  a  very 
unstable  substance,  decomposing  on  being  mixed  with  copper  oxide, 
powdered  glass  or  the  like.  It  is  immediately  broken  up  by  alkalies 
and  even  by  sodium  acetate  at  ordinary  temperatures,  nitrous  oxide 
and  water  being  formed  in  the  latter  case. 


PHOSPHORUS. 

131.  Phosphorus  does  not  occur  free  in  nature,  inasmuch  as  it 
combines  very  easily  with  oxygen.  Nevertheless  salts  of  phos- 
phoric acid  are  widely  distributed  and  occur  in  large  quantities. 
Tricalcium  phosphate,  Ca3(PO4)2,  phosphorite,  is  found  in  large 
deposits;  other  phosphates  which  are  frequently  met  with  are: 
apatite,  3Ca3(PO4)2  +  CaCl2  (or  CaF2);  waveUite,  4A1PO4+ 
2A1(OH)3  +  9H2O;  and  vivianite,  Fe3(P04)2  +  8H2O.  Phosphates 
are  also  found  to  a  small  extent  in  granite  and  volcanic  rocks; 
when  these  weather  the  phosphates  enter  the  soil.  About  0.1% 
of  phosphates  (calculated  as  P2O5)  is  present  in  soil  of  average 
fertility.  Bones  contain  a  considerable  proportion  of  tricalcium 
phosphate. 

Phosphorus  has  been  known  for  a  long  time.  In  1674  the 
alchemist  BRAND  discovered  it  at  Hamburg  by  evaporating  urine 
and  heating  the  residue  with  sand  in  an  earthen  retort.  SCHEELE 


200  INORGANIC  CHEMISTRY.  [§§  131- 

first  prepared  phosphorus  from  bones  by  a  process  which  in  the 
main  is  still  employed. 

For  the  manufacture  of  phosphorus  bones  or  phosphorite  are 
used.  By  treatment  with  sulphuric  acid  gypsum  and  monocalcium 
phosphate  are  formed.  As  the  latter  remains  soluble  while  the 
former  is  precipitated,  the  phosphate  solution  is  readily  poured  ofL 

Ca3(PO4)2+2H2SO4=2CaS04+CaH4(P04)2. 

SNSS,^  Gypsum. 

On  being  heated  this  phosphate  loses  water  and  is  converted  into 
calcium  metaphosphate : 

CaH4(P04)2  -  Ca(P03)2  +  2H2O. 

The  latter  [or  in  some  processes  the  original  phosphate  material] 
is  mixed  with  powdered  coal  and  sand  and  reduced  in  a  con- 
tinuously operated  electric  arc  furnace,  whereupon  phosphorus 
distils  over: 

2Ca(P03)2+2Si02  +  lGC=2CaSi03  +  10CO+4P. 

The  phosphorus  thus  obtained  is  black  because  of  adhering 
particles  of  charcoal,  and  contains  still  other  impurities.  It  can  be 
purified  by  distilling  with  steam  in  an  atmosphere  of  carbon  dioxide. 
Finally  the  phos'phorus  is  cast  into  sticks  and  in  this  form  enters 
the  market. 

132.  Physical  Properties. — At  ordinary  temperatures  phos- 
phorus is  a  crystalline  solid  of  a  light  yellow  color  and  having  a 
specific  gravity  of  1.83  at  10°.  When  cold  it  is  brittle  on  account 
of  its  crystalline  texture;  above  15°  it  becomes  soft  and  waxy  and 
melts  at  44.4°  to  a  yellow,  strongly  refractive  liquid.  Its  boiling- 
point  is  290°;  it  then  turns  to  a  colorless  vapor.  In  the  sunlight 
it  becomes  yellow  and  coated  with  an  opaque  pink  layer.  Phos- 
phorus is  insoluble  in  water,  slightly  soluble  in  alcohol  and  ether, 
but  easily  soluble  in  carbon  disulphide,  from  which  it  crystallizes 
in  crystals  of  the  regular  system. 

The  vapor  density  of  phosphorus  at  temperatures  between  515° 
and  1040°  is  4.58-4.50  (air  =  l).  Its  molecular  weight  is  there- 
fore 123.84.  Inasmuch  as  the  smallest  amount  of  phosphorus  that 
is  found  in  one  gram-molecule  of  any  of  the  numerous  phosphorus 
compounds  investigated  is  31  g.;  this  number  therefore  represents 


133.]  PHOSPHORUS.  201 

the  atomic  weight  of  the  element;  hence  the  molecule  must  consist 

123  84 
of        i     =  4  atoms.     Between  1500°  and  1700°  the  vapor  density 

decreases  considerably,  but  does  not  quite  reach  the  value  corre- 
sponding to  P2  molecules.  By  the  cryoscopic  method  also  (§  43) 
it  has  been  found  that  the  molecule  of  phosphorus  consists  of 
four  atoms  at  ordinary  temperatures. 

133.  Allotropic  Forms. — Ordinary,  or  yellow,  phosphorus  can 
be  transformed  by  heating  to  250°-300°  (in  absence  of  air,  neces- 
sarily) into  a  reddish-brown  powder,  red  phosphorus.  Iodine 
is  an  accelerator  of  this  process,  so  that  a  very  small  quantity  makes 
it  possible  for  this  transition  to  take  place  even  below  200°.  Red 
phosphorus  is  manufactured  on  a  large  scale  by  heating  the  yellow 
form  in  sealed  iron  cylinders  for  a  few  minutes  at  250°-300°. 
After  cooling,  the  product  is  treated  with  carbon  disulphide  and 
with  caustic  soda  in  order  to  extract  the  unchanged  yellow  phos- 
phorus. 

Red  phosphorus  is  considerably  different  in  its  properties  from 
the  yellow  form.  It  is  not  poisonous,  while  the  latter  is  very 
much  so.  Red  phosphorus  oxidizes  very  slowly  in  the  air; 
it  is  insoluble  in  carbon  disulphide.  Moreover  it  is  odorless, 
while  yellow  phosphorus  gives  off  a  peculiar  odor  because  of 
the  formation  of  ozone  (§  36).  When  heated  in  the  air,  it 
does  not  ignite  till  a  temperature  of  260°  is  reached.  On 
the  whole  it  is  chemically  much  less  active  than  yellow  phos- 
phorus. 

The  yejlow  modification  can  be  kept  for  an  unlimited  period 
at  ordinary  temperatures;  nevertheless  it  must  be  regarded  as  an 
unstable,  or,  rather,  metastable,  form,  just  like  "  detonating-gas," 
for  at  an  elevated  temperature  (the  higher,  the  faster),  both  pass 
over  into  more  stable  forms  with  the  evolution  of  heat.  The 
relative  stability  of  the  yellow  phosphorus  is  merely  a  consequence 
cf  the  extraordinarily  small  velocity  with  which  the  transfor- 
mation into  the  red  form  takes  place  at  ordinary  temperatures, 
notwithstanding  that  the  caloric  effect  of  the  change  is  consider- 
able, viz.,  32  Cal.  Yellow  phosphorus  can  be  regained  from 
the  red  form  by  distilling  the  latter  and  cooling  the  vapor  rapidly. 
It  is  a  general  rule  that  a  substance  which  is  capable  of  existing  in 
different  forms  appears  in  the  least  stable  one  first.  Now,  if  the 


202  INORGANIC    CHEMISTRY.  [§§134- 

temperature  is  low,  the  velocity  of  transformation  of  the  yellow 
form  is  so  small  that  it  seems  perfectly  stable. 

The  red  phosphorus  is  not  a  homogeneous  substance;  if  it  is  stirred 
up  with  water  some  violet-colored  particles  of  metallic  lustre  and  a 
specific  gravity  of  2.34  are  left  on  the  bottom.  HITTORF  prepared 
this  metallic  phosphorus  many  years  ago  by  heating  phosphorus  with 
lead  in  a  sealed  tube;  the  phosphorus  dissolves  in  molten  lead  and 
crystallizes  out  in  the  metallic  form  on  cooling.  The  specific  gravity  of 
the  red  phosphorus  varies  between  2.1  and  2.28,  according  to  the  duration 
and  temperature  of  the  heating.  It  is  probable  that  it  represents  a 
solid  solution  of  white  phosphorus  in  metallic  phosphorus.  The  ratio 
in  which  these  two  modifications  occur  in  red  phosphorus  changes  with 
the  temperature;  an  equilibrium  establishes  itself  between  them  for 
every  temperature.  At  about  500°  the  proportion  of  white  phosphorus 
reaches  a  minimum. 

Accordingly  only  the  white  and  the  metallic  phosphorus  are  to  be 
regarded  as  well  characterized  allotropic  modifications. 

134.  Chemical  Properties. — Phosphorus  has  a  great  affinity  for 
many  elements;    it  combines  directly  with   all   elements   except 
nitrogen  and  carbon,  the  combination  occurring  with  great  vigor 
in  many  cases,  e.g.  when  phosphorus  is  brought  in  contact  with 
sulphur  or  bromine.     Certain   compounds   of  the  metals    (phos- 
phides) are  known,  which  are   called  phosphor  bronzes  (§  199). 
Especially  characteristic  of  phosphorus  is  its  very  strong  affinity 
for  oxygen;    yellow  phosphorus  takes  fire  in  the  air  at  40°,  so 
that  contact  with  a  hot  glass  rod  is  sufficient  to  ignite  it.     The 
burning  is  accompanied  by  a  vigorous  evolution  of  light  and  heat, 
phosphorus  pentoxide,  1*265,  being  formed.     On  account  of  this 
strong  affinity  for  oxygen  phosphorus  is  a  powerful  reducing-agent. 
Sulphuric  acid,  when  warmed  with  it,  is  reduced  to  sulphur  dioxide; 
concentrated  nitric  acid  oxidizes  it  with  explosive  violence;  dilute 
acid  evolves  nitrous   fumes,   oxidizing  the  phosphorus   to  phos- 
phoric acid.     Many  metals  are  precipitated  by  phosphorus  from 
their   salts,    phosphides    being    formed    to    some    extent.     Silver 
nitrate,  for  instance,  gives  silver  and  silver  phosphide,  Ag3P,  with 
phosphorus;    on  warming  phosphorus  with  a  solution  of  copper 
sulphate  copper  phosphide,  Cu3P2,  is  deposited. 

135.  The  slow  oxidation  of  phosphorus  by  oxygen  at  ordinary 


135.]  PHOSPHORUS.  203 

temperatures  is  accompanied  by  the  emission  of  a  bluish  light. 
This  luminosity  of  phosphorus  is  very  plain  in  the  dark. 

This  phenomenon  is  due  to  various  circumstances,  some  of  which 
are  very  mysterious.  The  oxidation,  and  hence  the  luminosity,  is 
prevented  by  the  presence  of  traces  of  certain  substances,  such  as  hydro- 
carbons, ammonia,  etc.  Further,  the  luminosity  depends  on  the  tem- 
perature; below  10°  it  is  extremely  weak.  The  gas  pressure  has  a 
peculiar  influence;  at  ordinary  temperatures  phosphorus  does  not  emit 
light  in  pure  oxygen  of  atmospheric  pressure,  but  if  the  pressure  is 
reduced,  a  point  is  reached  at  which  luminosity  commences;  this  is  at 
666  mm.. for  15°,  and  at  760  nini.  for  19.2°.  The  oxidation  is  there- 
fore more  vigorous  in  dilute  oxygen  (i.e.  oxygen  mixed  with  another 
gas,  such  as  nitrogen)  than  in  concentrated.  VAN  MARUM  observed 
as  early  as  1798  that  a  piece  of  phosphorus  laid  on  wadding  (which 
serves  as  a  poor  conductor  of  heat)  in  a  closed  vessel  shines  the  more 
brightly  as  the  oxygen  is  pumped  out,  and  may  even  take  fire  hi  very 
dilute  gas. 

The  fact  that  oxidations  are  more  energetic  under  reduced  oxygen- 
pressure  has  been  observed  in  many  other  cases.  See  §  137. 


Detection  of  Phosphorus. — Poisonings  by  yellow  phosphorus 
occur  now  and  then.  In  order  to  detect  it  in  such  cases,  use  is 
made  of  its  luminosity.  For  this  purpose  the  contents  of  the 
stomach,  which  are  to  be  tested  for  phosphorus,  are  diluted  with 
water  in  a  distilling-flask,  connected  with  a  condenser  by  a  tube 
doubly  bent  at  right  angles.  On  heating  the  flask  water  distils 
over  with  a  little  phosphorus  vapor;  if  the  whole  apparatus  is 
placed  in  a  dark  room,  a  luminous  ring  is  noticed  during  this 
distillation  at  the  place  where  the  steam  is  condensed,  i.e.  where 
the  phosphorus  vapor  comes  in  contact  with  air  in  the  condenser. 
The  distillate  contains  phosphoric  acil  (MITSCHERLICH'S  test). 

Use.  —  Phosphorus  is  used  chiefly  for  the  manufacture  of 
m  a  t  c  h  e  s.  The  matches  in  use  to-day  may  be'classed  as  safety  matches 
and  "strike  any  where"  matches,  which  latter  maybe  of  the  (a)  parlor 
single-dip  type,  or  (6)  double-dip  with  combustible  bulb  or  (c)  double- 
dip  with  safety  bulb.  In  the  Swedish  safety  matches  the  head  consists 
chiefly  of  a  mixture  of  potassium  chlorate  and  antimony  sulphide. 
They  are  lighted  by  striking  them  on  a  surface  coated  with  red  phos- 


204  INORGANIC  CHEMISTRY.  [§§  135- 

phorus.  On  account  of  their  requiring  a  special  ignition  surface  and 
their  quick  burning  and  dropping  they  are  not  as  popular  in  America 
as  the  strike-anywhere  matches.  Of  the  latter  the  double-dip  safety 
bulb  matches  are  best  because  they  can  be  ignited  by  friction  of  the 
tip  only,  not  by  side  friction. 

Phosphorus,  in  some  form,  is  used  in  the  production  of  all  matches. 
Yellow  phosphorus  is  generally  utilized  in  the  production  of  American 
matches,  but  since  it  is  very  injurious  to  the  health  of  the  workmen, 
most  European  countries  forbid  its  use  and  employ  non-poisonous 
substitutes.  The  red  phosphorus  and  phosphorus  "  sesqui-sulphide,  " 
P4S3,  are  the  chief  non-poisonous  substitutes  for  yellow  phosphorus 
and  are  to  be  utilized  in  American  strike-anywhere  matches  as  fast 
as  practicable.  Compare  §  133. 

Compounds  of  Phosphorus  and  Hydrogen. 

There  are  three  compounds  of  phosphorus  and  hydrogen  known  : 

(1)  gaseous  hydrogen  phosphide,   PH3   (also   called   phosphine); 

(2)  liquid  hydrogen  phosphide,  P2H4;    and   (3)    solid  hydrogen 
phosphide,  (P2H)6. 

HYDROGEN   PHOSPHIDE.     PHOSPHINE,  PH3. 

136.  This  compound  can  be  prepared  from  the  elements  by 
bringing  phosphorus  together  with  zinc  and  dilute  sulphuric  acid, 
i.e.,  with  nascent  hydrogen;  when  thus  prepared,  it  is  mixed  with 
a  large  quantity  of  hydrogen. 

The  generation  of  hydrogen  phosphide  by  heating  phosphorous 
and  hypophosphorous  acids  is  another  example  of  its  formation  by 
the  direct  decomposition  of  phosphorus  compounds: 


Phosphorous  Phosphoric 

acid.  acid. 

The  ordinary  method  of  preparation  is  by  the  action  of  phos- 
phorus on  caustic  potash: 

P4+3KOH+3H20  =  PH3+3H2KP02. 

Pot.  hypophos- 

The  reaction  is  really  more  complicated  than  this  equation  indicates, 
for  in  addition  hydrogen,  P2H4  and  other  substances  are  formed.  (See 
also  1  144.) 


136.] 


HYDROGEN  PHOSPHIDE. 


205 


By  reason  of  the  presence  of  gaseous  P2H4,  which  is  self-inflammable, 
each  bubble  of  gas  ignites  as  it  breakes  into  the  air,  forming  usually  a 
smoky  ring  of  phosphorus  pentoxide  (Fig.  37).  On  account  of  this 
inflammability  the  vessel  in  which  the  gas  is  generated  from  phosphorus 
and  caustic  potash  must  be  as  full  of  liquid  as  possible.  Moreover,  the 
delivery-tube  (preferably  with  a  wide  mouth)  must  open  in  warm  water, 
in  order  that  it  may  not  become  clogged  with  particles  of  phosphorus 
carried  over.  By  passing  the  gas  through  hydrochloric  acid  or  alcohol, 
the  hydrogen  phosphide  is  freed  from  P2H4  and  is  then  no  longer  spon- 
taneously combustible. 


FIG.  37. — PREPARATION  OF  HYDROGEN  PHOSPHIDE. 

According  to  SENDERENS  phosphine  can  be  very  advantageously 
prepared  by  heating  red  phosphorus  at  240°-250°  in  a  current  of  steam. 
Perhaps  phosphorus  or  metaphosphorous  acid  is  formed  primarily. 

No  method  of  producing  hydrogen  phosphide  by  the  action 
of  hydrogen  on  phosphorus  compounds  is  known;  however,  we 
have  one  by  the  interaction  of  hydrogen  compounds  and  phos- 
phorus compounds.  Calcium  phosphide,  when  decomposed  by 
water  or  dilute  hydrochloric  acid,  forms  hydrogen  phosphide: 

Ga3P2  +  6HC1  =  3CaCl2  +  2PH3. 
The  phosphides  of  zinc,  iron,  tin,  and  magnesium  are  decom- 


206  INORGANIC  CHEMISTRY.  [§§136- 

posed  by  dilute  acids  with  the  formation  of  hydrogen  phos- 
phide. 

The  perfectly  pure  gas  is  now  best  obtained  by  condensa- 
tion and  subsequent  fractionation  at  a  very  low  temperature,  as 
was  described  in  §  29. 

Physical  Properties. — Hydrogen  phosphide,  or  phosphuretted 
hydrogen,  PH3,  is  a  gas  at  ordinary  temperatures;  it  becomes 
liquid  at  —85°  and  solid  at  —133.5°.  It  has  a  peculiar  disagree- 
able odor,  that  reminds  one  of  spoiled  fish.  It  is  slightly  soluble 
in  water,  more  so  in  alcohol.  Sp.  g.=  17  (0=  16).  1  liter  weighs 
1,5293  g.  at  O°  and  760  mm.  pressure. 

137.  Chemical  Properties. — Hydrogen  phosphide  is  very  poison- 
ous; it  burns  very  easily,  yielding  phosphoric  acid.  In  the 
presence  of  oxygen  of  ordinary  pressure  it  remains  unchanged; 
if,  however,  the  pressure  is  diminished,  an  explosion  results.  This 
conduct  reminds  one  of  phosphorus,  which  is  luminous  (because 
of  oxidation)  only  below  a  certain  limit  of  pressure  (§  135). 

The  combustion  of  hydrogen  phosphide  may  be  expressed  by 
the  equation: 

2PH3  +402  =  P205  +  3H2O. 

Accordingly  the  reaction  would  be  hexamolecular  (§  50).  VAN 
DER  STADT  demonstrated  by  a  method,  similar  to  that  referred 
to  in  §  51,  that  the  first  stage  of  the  reaction  is  bimolecular  and 
corresponds  very  closely  to  the  following  equation: 

PH3  +  02=H2+P02H, 

Metaphos- 
phorous  acid. 

if  the  gases  slowly  diffuse  into  each  other  in  a  diluted  condition. 

In  general,  experience  has  taught  that  the  mechanism  of  a 
reaction  is  decidedly  simple  and  that  chemical  processes  are  almost 
always  mono-  or  bimolecular.  Accordingly,  when  the  quantita- 
tive course  of  a  reaction  is  represented  by  an  equation  indicating 
the  participation  of  several  molecules,  it  is  probable  that  several 
intermediate  reactions  are  involved. 

Hydrogen  phosphide  can  unite  with  halogen-hydrogen  acids 
directly  to  form  compounds  of  the  type  PH4X  (X=haolgen),  in 
analogy  with  ammonia.  The  best  known  of  these  compounds 


137.]  HYDROGEN  PHOSPHIDE.  207 

is  PH4I,  phosphonium  iodide,  a  colorless,  well-crystallized  com- 
pound, which  is  formed  when  dry  hydrogen  phosphide  and  hydro- 
gen iodide  are  mixed.  In  contact  with  water  it  breaks  up  into 
PH3  and  HI;  the  former  escapes  as  a  gas,  while  the  latter  remains 
dissolved  in  water. 

Phosphonium  iodide  is  very  unstable.  This  is  even  more  the 
case  with  phosphonium  bromide,  which  is  also  a  solid,  but  is 
completely  dissociated  into  the  two  hydrogen  compounds,  PH3 
and  HBr,  as  low  as  30°.  Phosphonium  chloride  is  dissociated 
even  at  ordinary  temperatures  and  pressures  and  can  only  exist 
below  14°  or  under  more  than  20  atm.  pressure.  Considering 
these  properties,  it  is  not  surprising  that  phosphonium,  PH4,— 
like  ammonium, — should  be  impossible  to  isolate.  No  other  acids 
except  those  mentioned  unite  with  hydrogen  phosphide.  The 
general  behavior  of  the  latter  thus  shows  that  it  is  very  much  less 
basic  than  ammonia. 

It  is  for  this  reason  that  PH4I  is  decomposed  by  water.  The  weak 
basic  character  of  phosphonium  hydroxide  allows  the  iodide  to  be 
hydrolyzed  into  PH4OH  and  HI,  whereupon  the  PH4OH  breaks  up 
forthwith  into  PH3  and  H2O. 

Hydrogen  phosphide  possesses  reducing  properties.  From 
solutions  of  silver  nitrate  of  copper  sulphate  it  precipitates  a 
mixture  of  metal  and  phosphide;  this  property  can  be  made  use 
of  to  separate  the  gas  from  other  gases.  When  mixed  with  chlo- 
rine it  explodes  vigorously,  forming  hydrochloric  acid  and  phos- 
phorus trichloride. 

The  composition  of  hydrogen  phosphide  was  determined  by 
passing  a  known  volume  over  copper  turnings  in  a  heated  tube. 
The  copper  combines  with  the  phosphorus,  so  that  the  increase 
in  weight  of  the  tube  shows  the  proportion  of  .phosphorus.  The 
escaping  hydrogen  is  collected  and  measured.  In  the  experi- 
ment it  was  found  that  hydrogen  phosphide  contains  91.2% 
phosphorus  and  8.8%  hydrogen.  The  specific  gravity  for  O=16 
was  found  to  be  17;  the  molecular  weight  is  therefore  34.  In 
34  parts  there  are,  according  to  the  above  composition,  31  parts 
of  phosphorus  and  3  parts  of  hydrogen;  the  composition  of  the 
compound  is  therefore  expressed  by  PH3.  This  agrees  with  the 
results  of  the  decomposition  of  hydrogen  phosphide  by  induction 


208  INORGANIC   CHEMISTRY.  [§§138- 

sparks  or  by  the  electric  arc;  1  vol.  PH3  yields  H  vols.  hydrogen 
and  also  amorphous  phosphorus,  which  is  deposited  on  the  sides 
of  the  tube  and  on  the  platinum  wires  (or  carbons). 


2PH3=3H2+2P. 

2  vols.      3  vols.    solid. 


Liquid  Hydrogen  Phosphide,  P2H4. 

138.  In  certain  cases  this  compound  is  formed  as  a  by-product  in 
the  preparation  of  phosphine.  Especially  is  it  formed  in  the  decom- 
position of  calcium  phosphide  with  water.  It  also  results  from  the 
>xidation  of  phosphine  by  various  substances,  for  instance,  nitrogen 
dioxide,  by  which  ordinary  hydrogen  phosphide  can  be  made  spon- 
taneously inflammable.  The  mixture  of  PH3  and  P2H4  can  be  separated 
by  passing  it  through  a  well-cooled  tube;  the  latter  substance  con- 
denses to-  a  colorless  liquid,  which  boils  at  57°-58°  (under  735  mm.) 
and  has  a  specific  gravity  of  1.01.  It  is  easily  decomposable  and  can- 
not be  preserved,  because  it  rapidly  changes  to  the  gaseous  and  the 
solid  hydrogen  phosphides.  The  same  decomposition  is  also  effected 
by  hydrochloric  acid.  It  must  be  condensed  in  the  dark,  as  sunlight 
aids  decomposition.  The  empirical  composition  is  indicated  by  the 
formula  PH2;  but  since  phosphorus  is  trivalent,  we  take  P2H4,  i.e. 
H2P-PH2,  as  the  formula  of  the  molecule.  Liquid  hydrogen  phosphide 
thus  becomes  analogous  to  hydrazine. 


Solid  Hydrogen  Phosphide,  P12H6. 

139.  This  substance  is  formed  by  the  decomposition  of  the  preceding 
one, — especially  easily  when  phosphine  charged  with  P2H4  vapor,  as  it 
is  evolved  from  the  decomposition  of  Ca3P2  with  water,  is  led  over  granu- 
lated calciumchloride.  The  solid  hydrogen  phosphide  separates  out  as  a 
bright  yellow  powder  of  the  empirical  formula  P2H,  whose  molecular  weight 
has  been  determined  crysocopically  to  be  P12H6  (using  yellow  phosphorus 
as  a  solvent).  On  being  heated  to  about  200°  it  breaks  up  into  phos- 
phine and  a  new  solid  hydrogen  phosphide  of  the  empirical  formula 
P9H2.  This  latter  compound  is  also  formed  when  P12H6  is  treated  with 
liquid  ammonia.  When  heated  in  the  air  P12H8  catches  fire  at  160°. 
It  is  insoluble  in  water. 


140.]  PHOSPHORUS  TRICHLORIDE.  209 

Halogen  Compounds  of  Phosphorus. 

Phosphorus  unites  with  all  four  of  the  halogens  to  form  com- 
pounds of  the  types  PXs  and  PX5;  the  most  important  are  the 
chlorides. 

PHOSPHORUS  TRICHLORIDE,  PC13. 

140.  This  compound  can  only  be  obtained  by  direct  combina- 
tion of  the  elements.  In  preparing  it  a  rapid  current  of  dry 
chlorine  is  led  over  phosphorus  in  a  retort.  The  phosphorus  burns 
with  a  pale  yellow  fiame  and  a  mixture  of  trichloride  and  penta- 
chloride  distils  over  into  the  receiver,  which  is  kept  cold.  A  little 
phosphorus  is  added  to  the  distillate  in  order  to  convert  the  pen- 
tachloride  to  trichloride,  and  the  liquid  is  redistilled.  An  easier 
method  is  to  introduce  phosphorus  into  a  flask  with  some  phos- 
phorus trichloride  and  lead  chlorine  into  the  mixture. 

Physical  Properties. — Phosphorus  trichloride  is  a  colorless 
liquid  of  a  very  pungent  odor;  it  boils  at  76°  and  remains  liquid 
as  low  as  -115°.  Sp.  g.  =  1.6129  at  0°. 

Chemical  Properties. — Water  decomposes  it  very  rapidly  with 
the  formation  of  hydrochloric  and  phosphorous  acids: 

PC13  +  3H20  =  HaPO3 + 3HC1. 

It  is  because  of  this  decomposition  that  it  fumes  in  moist  air. 
Continued  treatment  with  chlorine  converts  it  into  the  penta- 
chloride. 


210  INORGANIC  CHEMISTRY.  [§§  141- 


PHOSPHORUS  PENTACHLORIDE,  PC16. 

141.  This  substance  is  prepared  by  passing  chlorine  over  phos* 
phorus  trichloride.  Fine  light-yellow  crystals  at  once  appear  and 
the  entire  mass  finally  becomes  solid,  indicating  that  all  is  con- 
verted into  penta  chloride.  This  compound  fumes  strongly  in 
moist  air,  being  immediately  decomposed  by  watei  with  the  forma- 
tion of  hydrochloric  and  phosphoric  acids.  When  heated  it  sub- 
limes without  melting.  In  the  transition  to  the  gaseous  state  it 
breaks  up  at  a  rather  low  temperature  into  the  trichloride  and 
chlorine;  this  dissociation  is  complete  at  300°,  for  at  that  point 
the  vapor  density  is  just  half  of  that  calculated  for  the  penta- 
chloride.  The  vapor,  which  at  moderately  low  temperatures  is 
almost  colorless,  takes  on  the  yellow  color  of  chlorine  for  the  above 
reason,  as  the  temperature  rises.  The  dissociation  products, 
phosphorus  trichloride  and  chlorine  can  be  separated  by  diffusion. 
Phosphorus  pentachloride  evaporates  in  an  atmosphere  of  the  tri- 
chloride with  almost  no  dissociation  (cf.  §  51). 

By  the  addition  of  a  little  water  it  is  converted  into  phosphorus 
oxy  chloride: 

PC15+H20  =  POC13  +  2HC1. 

With  more  water  phosphoric  and  hydrochloric  acids  are  produced. 
Phosphorus  pentachloride  is  used  in  organic  chemistry  to 
replace  hydroxyl  groups  with  chlorine.  In  inorganic  chemistry, 
it  can  also  be  employed  for  the  same  purpose;  thus  sulphuric  acid 
reacts  with  it  in  the  following  manner  (§  89)  : 


PHOSPHORUS  OXYCHLORIDE,  POC13. 

142.  The  best  method  of  preparing  this  compound  is  by  the 
oxidation  of  the  trichloride  with  sodium  chlorate: 

3PC13  +  NaClO8  =  3POC13  +  NaCl. 
In  order  to  moderate  the  great  vigor  of  this  reaction  sodium  chlorate 


144.]  COMPOUNDS  OF  PHOSPHOROUS.  211 

is  placed  under  phosphorus  oxy chloride  and  the  trichloride  is 
then  added  slowly  by  means  of  a  dropping-funnel. 

Phosphorus  oxychloride  is  a  colorless,  mobile  liquid,  that  boils 
at  107.2°  and,  when  solid,  melts  at  -1.5°.  Sp.  g.  =1.7118  at  0°. 
In  the  presence  of  water,  with  which  it  is  not  miscible,  it  slowly 
changes  to  phosphoric  and  hydrochloric  acids: 

POC13 + 3H20  =  H3P04+ 3HC1. 

THE  COMPOUNDS  OF  PHOSPHORUS  WITH  THE  OTHER  HALOGENS. 

143.  These  are  very  analogous  to  the  chlorine  derivatives.  They  are 
likewise  prepared  by  direct  synthesis  from  the  elements.  Inasmuch  as 
the  reaction  is  very  vigorous,  it  has  to  be  moderated  by  dissolving  the 
phosphorus  and  the  halogen  separately  in  carbon  disulphide,  slowly 
adding  one  to  the  other  and  then  distilling  off  the  solvent.  The  fluorides 
have  special  methods  of  preparation.  All  these  compounds  are  broken 
up  by  water  like  the  corresponding  chlorides,  the  fluorides,  however, 
quite  slowly. 

The  composition  of  these  compounds  can  be  ascertained  in  the 
following  way:  on  being  decomposed  by  water  they  yield  phos- 
phoric or  phosphorous  acid  and  a  halogen  acid,  so  that  the  quan- 
tities of  phosphorus  and  halogen  present  can  be  found  by  deter- 
mining the  amounts  of  these  acids.  Moreover,  the  molecular 
weight  can  be  obtained  by  measuring  the  vapor  density,  though 
it  must  be  borne  in  mind,  however,  that  compounds  of  the  type 
PX5  are  usually  dissociated  in  the  gaseous  state. 

Oxygen  Compounds  of  Phosphorus. 

144.  Three  compounds  of  this  class  are  known:  phosphorus 
trioxide,  P20s;  phosphorus  tetroxide,  P204,'  and  phosphorus  pent' 
oxide,  or  phosphoric  anhydride,  P2O5.  Only  the  last  is  of  any  great 
importance. 


Phosphorus  Trioxide,  P203. 

This  compound  is  produced  when  phosphorus  burns  in  a  slow  cur- 
rent of  dry  air  in  a  tube.  The  principal  product  is  phosphorus  pent- 
oxide,  which  can  be  collected  by  a  wad  of  glass  fibers.  The  phosphorus 
trioxide  passes  through  as  a  vapor  and  is  condensed  in  a  well-cooled 


212  INORGANIC  CHEMISTRY.  [§§144- 

tube.  It  is  a  white  waxy  substance  when  thus  formed,  but  it  can  also 
be  obtained  in  crystals;  the  latter  melt  at  22.5°  and  boil  at  173.1°  (in 
a  nitrogen  atmosphere).  The  vapor  density  has  been  found  to  be  109.7, 
while  that  calculated  for  P406  is  110.  On  being  heated  to  440°  it  is 
decomposed  into  red  phosphorus  and  phosphorus  tetroxide.  It  turns 
yellow  in  the  light,  which  explains  the  fact  that  *Pfosphorus  pentoxide 
sometimes  takes  on  a  yellow  color.  It  dissolves  slowly  in  cold  water 
forming  phosphorous  acid;  with  hot  water  it  produces  red  phosphorus, 
self-inflammable  hydrogen  phosphide  and  phosphoric  acid  in  a  vigor- 
ous reaction.  When  heated  to  50°-60°  in  the  air  it  takes  fire  and  burns 
to  the  pentoxide. 

Phosphorus  Tetroxide,  P204, 

is  obtained  from  the  P203  compound,  as  was  stated  above.  It  forms 
colorless  glistening  crystals,  that  break  up  in  water  into  phosphorous 
and  phosphoric  acids.  In  this  respect  its  conduct  is  analogous  to  that 
of  nitrogen  tetroxide,  which  yields  nitrous  and  nitric  acids  with  water. 

PHOSPHORUS  PENTOXIDE,  P205. 

This  compound  is  the  product  of  the  combustion  of  phosphorus 
in  oxygen  or  an  excess  of  dry  air.  It  forms  a  white,  voluminous, 
snow-like  mass,  that  takes  up  water  rapidly  to  produce  phosphoric 
acid.  It  is  the  most  powerful  desiccating-agent  known.  MORLEY 
ascertained  that  it  dries  the  air  down  to  1  mg.  water  vapor  in 
40,000  1.  air.  It  exists  in  two  modifications,  both  of  which  are 
formed  simultaneously  in  the  above  process.  The  one  is  crystal- 
line, subliming  at  250°;  the  other  amorphous  and  not  volatile 
below  red  heat;  the  vapor  condenses  crystalline.  When  heated 
above  250°  the  crystalline  modification  passes  over  into  the 
amorphous  form. 

Heating  with  charcoal  reduces  it  to  phosphorus. 

The  vapor  density  of  phosphoric  anhydride  at  bright  redness 
was  found  to  correspond  to  the  formula  (P2O5)2. 

Acids  of  Phosphorus. 

145.  Only  two  of  the  above  described  oxides  of  phosphorus, 
viz.  P2O3  and  P205,  form  corresponding  acids;  these  oxides  can 
unite  with  different  amounts  of  water  to  form  acids.  From  P2Os 
we  have: 


145.]  ACIDS  OF  PHOSPHOROUS.  213 


P20s  +   H2O  =  2HP03,  metaphosphoric  acid, 

H4P2O7,  pyro  phosphoric  acid,  and 
=  2H3PO4,  orthophosphoric  acid. 


From  the  other^xide  two  acids  can  be  derived:  metaphos- 
phorous  acid,  HPO^phd  phosphorous  acid,  H3PO3.  Besides  these 
there  are  two  acids  of  phosphorus,  whose  anhydrides  are  unknown. 
viz.  hypophosphorous  acid,  H3PO2;  and  hypophosphoric  acid,  H4P2O6. 

The  relation  between  ortho-,  meta-,  and  pyrophosphoric  acids  can 
be  shown  in  another  way,  which  leads  us  to  make  some  general  obser- 
vations. It  was  remarked  in  §  141  that  phosphorus  pentachloride  is 
transformed  by  water  into  phosphoric  and  hydrochloric  acids.  The 
action  of  water  on  the  pentachloride  may  be  regarded  as  .consisting 
first  of  a  substitution  of  all  five  chlorine  atoms  by  hydroxyl: 

P|C15  +  5H|OH  =5HC1  +  P(OH)5. 

This  compound,  which  would  strictly  be  regarded  as  orthophos- 
phoric acid,  is  unknown;  a  molecule  of  water  is  at  once  split  off,  form- 
ing the  ordinary  phosphoric  acid,  H3PO4,  which  we  are  accustomed 
to  call  orthophosphoric  acid.  In  a  similar  way  the  metaphosphoric 
acid  can  be  derived  from  the  acid  P(OH)5  by  the  splitting  off  of  two 
molecules  of  water: 


C)H 

IQH  OH  IOH 

P  O  H  -»  OP^OH;          P  OJH  ->  O2P-OH; 

OH  ^OH  |OH       Metaph9sphoric 

O  H        Orthophos-  O  H 

phoric  acid. 

while  the  pyrophosphoric  acid  can  be  regarded  as  2P(OH)5-3H20: 

0|H  H]<3 

IOH  H_0|  .OIL  VOH 

P  O  H  H  O  P  ->  OP^-OH        OP—  OH 

OH  HO  \o/ 

IOH  HIO 

J  -  !  Pyrophosphoric  acid. 

Orthophosphoric  acid  can  also  be   derived  from  phosphorus  oxy- 
chloride  : 

OP(OH)3. 


214  INORGANIC  CHEMISTRY.  [§§  145- 

This  way  of  looking  at  them  makes  plain  not  only  the  connection 
between  the  different  acids,  but  also  their  structural  formulae.  The 
same  method  can  be  applied  to  many  other  cases.  As  an  example 
we  may  select  the  per-iodic  acids.  In  §  62  only  one  was  mentioned. 
There  are  salts,  however,  of  various  per-iodic  acjdg,  e.g.  MIO4,  M3IO5, 
M5I06,  etc.  These  can  be  derived  from  a  hypothetical  acid  I(OH)7 
in  which  iodine  is  joined  to  as  many  hydroxyls  as  correspond  to  its 
maximum  valence.  M5IO6  would  come  from  I(OH)7  — 1H2O;  M3IO5 
from  I(OH)7-2H2O;  and  MIOd  from  I(OH)7-3H2O. 

ORTHOPHOSPHORIC  ACID,  H3P04. 

146.  This  acid  can  be  obtained  by  direct  synthesis  from  its 
elements;  phosphorus  burns  to  the  pentoxide  and  the  latter  yields 
the  acid  on  dissolving  in  water.  Its  formation  by  the  action  of 
nitric  acid  on  phosphorus  was  mentioned  in  §  134.  It  can  also  be 
obtained  by  the  oxidation  of  compounds  containing  phosphorus 
and  hydrogen;  phosphine  and  the  lower  acids  of  phosphorus  are 
oxidized  to  phosphoric  acid. 

Ordinarily  this  acid  is  prepared  by  the  oxidation  of  phosphorus 
with  nitric  acid  or  by  liberating  it  from  its  salts,  particularly  the 
calcium  salt,  Ca3(PO4)2-  The  latter  is  stirred  into  the  theoretical 
amount  of  dilute  sulphuric  acid,  forming  calcium  sulphate,  which 
is  only  slightly  soluble  in  water,  and  phosphoric  acid,  which  goes 
into  solution.  On  evaporating  this  solution  the  acid  remains. 

At  ordinary  temperatures  orthophosphoric  acid  is  a  crystalline 
solid.  It  melts  at  38.6°,  is  odorless  and  extremely  soluble  in  water, 
forming  a  strongly  acid  solution. 

It  has  the  character  of  a  strong  acid;  however,  it  is  consider- 
ably less  ionized  than  hydrochloric  acid;  a  solution  of  1  mole 
phosphoric  acid  in  10  1.  water  contains  about  one-fourth  as  many 
hydrogen  ions  as  hydrochloric  acid  of  the  same  molecular  con- 
centration. It  is  ionized  chiefly  into  H'  and  H^PO/.  It  generates 
hydrogen  with  metals,  all  three  hydrogen  atoms  being  replaceable 
by  metallic  atoms;  it  is  therefore  tribasic.  Three  classes  of  salts 
are  possible  and  known  to  exist;  these  are  the  primary,  secondary 
and  tertiary  salts.  Of  the  alkali  salts  all  three  kinds  are  soluble; 
of  the  alkaline  earth  salts  only  the  primary,  the  tertiary  and 
secondary  being  insoluble.  The  other  phosphates  are  insoluble 
in  water  but  are  dissolved  by  mineral  acids. 


146.]  ORTHOPHOSPHORIC  ACID.  215 

This  latter  property  is  due  to  the  fact  that  phosphoric  acid  is 
a  weaker  acid  than  the  strong  mineral  acids,  hydrochloric,  nitric 
and  sulphuric.  On  treating  an  insoluble  phosphate  with  one  of 
these  acids,  e.g.  'hydrochloric,  undissociated  molecules  of  phos- 
phoric acid  are  formed  in  the  liquid;  the  more  hydrochloric  acid, 
the  more  the  association,  since  the  hydrochloric  acid  reduces  the 
ionization  of  phosphoric  acid.  H2PO4'  and  H'  ions  thus  disappear 
and,  in  case  enough  hydrochloric  acid  is  added,  the  concentration 
of  the  H2P04'  ions  remaining  will  not  be  great  enough  together 
with  that  of  the  metal  ions  present  to  reach  the  value  of  the  solu- 
bility product;  hence  all  the  phosphate  must  dissolve  (§  73). 

For  the  same  reason,  as  a  general  rule,  salts  that  are  insoluble 
in  water  will  only  dissolve  in  acids  that  are  stronger  than  the  acid 
of  the  salt.  The  only  exception  to  this  is  the  case  when  the  value 
of  the  solubility  product  of  the  insoluble  salt  is  very  small,  examples. 
of  which  we  have  seen  in  certain  sulphides  (§  73). 

When  heated  to  213°  orthophosphoric  acid  gives  off  water, 
forming  mainly  the  pyro-acid  but  also  a  little  meta-acid  through- 
out the  reaction.  The  pyro-acid  on  the  other  hand  is  converted 
by  further  heating  into  the  meta-acid. 

With  silver  nitrate  orthophosphates  give  a  yellow  precipi- 
tate of  silver  phosphate,  Ag3PO4,  soluble  in  nitric  acid  and  ammo- 
nia. In  the  case  of  a  primary  or  secondary  phosphate,  the  pre- 
cipitation is  not  complete,  since  nitric  acid  is  liberated  in  the 
reaction: 


Na2HP04  +3AgN03  =  AggPO*  +2NaNO3  +HN03, 
or,  expressed  in  ions: 

HP04"  +  3Ag-  <±  Ag3P04+H'. 

If,  however,  an  excess  of  sodium  acetate  is  added,  the  precipi- 
tation is  practically  complete. 

The  reason  for  this  is  obvious.  By  the  addition  of  acetate  the 
acetic  anions  C2H302'  are  forced  to  combine  with  the  H'  ions,  for 
acetic  acid  is  only  very  slightly  ionized  and  its  ionization  is,  more- 
over, considerably  lessened  by  the  excess  of  sodium  acetate.  The 
result  is  that  in  the  equilibrium  HPO4"  +  3Ag'  <=±  Ag3P04+H* 
the  H'  ions  are  removed.  The  inverse  reaction  <—  is  then  no  longer 


216  INORGANIC  CHEMISTRY.  [§§  146- 

possible,  and  the  direct  reaction  — >  must  therefore  become  com- 
plete, or  in  other  words,  all  the  phosphoric  acid  is  precipitated 
as  silver  phosphate. 

It  was  stated  above  that  the  alkali  salts  of  phosphoric  acid 
are  soluble  in  water.  These  aqueous  solutions  differ  markedly 
in  reaction.  The  solution  of  a  primary  salt,  KH2P04,  is  acid, 
that  of  a  secondary  salt  feebly  alkaline,  and  that  of  a  tertiary 
salt  strongly  alkaline.  The  cause  of  this  variation  must  be  more 
fully  explained. 

The  acid  reaction  of  a  salt  such  as  KH2PO4  must  be  attributed 
to  the  fact  that  its  anion,  H2PO4'  (analogous  to  the  anion  HSO4'), 
is  capable  of  splitting  up  into  the  ions  H'  and  HPO4",  the  former 
producing  the  acid  reaction.  The  feebly  alkaline  reaction  of  a 
salt  like  K2HPO4  is  accounted  for  by  hydrolysis  (§  66).  Such 
a  salt  is  extensively  ionized  in  dilute  solution  into  2K'  and  HPO4". 
However,  while  H3P04  is  rather  highly  dissociated  (into  H*  and 
H2PO4'),  H2PO4'  is  but  slightly  ionized  into  H*  and  HPO4". 
In  this  case  H2P04'  behaves  as  a  weak  acid.  Hence,  if  there  is 
a  large  proportion  of  HPO4"  ions  in  a  solution,  they  will  tend 
to  unite  with  H'  ions,  because  the  system  H*  +  HPO4":^±H2PO4' 
is  only  in  equilibrium  when  the  right-hand  side  preponderates. 
The  necessary  H'  ions  are  supplied  by  the  water,  which  is  split 
up  to  a  very  slight  extent  into  H'  and  OH'.  But  when  the  H* 
ions  unite  with  HP04"  ions  we  have  a  surplus  of  OH'  ions  in 
the  solution  and  the  latter  takes  on  an  alkaline  reaction.  Entirely 
analogous  is  the  explanation  of  the  strongly  alkaline  reaction  of 
the  tertiary  phosphates,  such  as  K3PO4.  Their  aqueous  solutions 
contain  the  ions  PO4'",  which  have  a  still  stronger  tendency  to 
unite  with  H'  ions  than  the  HPO4"  ions.  The  P04'"  ion,  there- 
fore, causes  the  presence  of  an  even  larger  proportion  of  OH'  ions, 
not  compensated  by  H'  ions,  so  that  the  result  is  a  strongly  alkaline 
reaction. 

Phosphoric  acid  is  precipitated  from  an  ammoniacal  solution 
by  a  magnesium  salt  as  white  crystalline  ammonium  magnesium 
phosphate,  NH4MgPO4+6H20.  Another  very  characteristic  test 
for  phosphoric  acid  is  that  in  nitric  acid  solution  a  finely  crystal- 
line, yellow  precipitate  is  produced  by  ammonium  molybdate, 
especially  on  warming.  This  precipitate  has  approximately  the 
composition  14Mo03+(NH4)3P04+4H20,  i.e.  it  is  an  ammo- 


148.]  PYROPHOSPHORIC  ACID.  217 

mum  phospho-molybdate.  Precipitation  in  acid  solution  is  of 
great  advantage  here;  since  most  of  the  phosphates  are  soluble 
only  in  acids. 

PYROPHOSPHORIC  ACID,  H4P207. 

147.  One  method  of  producing  this  acid  was  given  in  the  pre- 
ceding paragraph.     In  preparing  it,  it  is  more  practicable,  however, 
to  heat  the  secondary  sodium  phosphate   (the  ordinary  sodium 
phosphate  of  commerce),  because  in  this  case  only  one  molecule 
of  water  can  be  driven  off  from  two  molecules  of  the  salt: 

2Na2HPO4= H2O + Na4P2O7. 

After  being  heated  the  sodium  pyrophosphate  is  dissolved  in  water 
and  lead  acetate  is  added  to  precipitate  lead  pyrophosphate,  which 
is  then  decomposed  with  hydrogen  sulphide. 

Pyrophosphoric  acid  can  be  obtained  from  its  solution  as  a 
colorless  vitreous  mass  by  .evaporation  in  a  vacuum  at  a  low  tem- 
perature. When  dissolved  in  water  of  ordinary  temperature,  the 
acid  remains  unchanged  for  quite  a  while ;  on  warming  this  solu- 
tion, especially  after  the  addition  of  a  little  mineral  acid,  it  is 
converted  in  a  few  hours  into  ortho-acid  (§  145). 

All  four  hydrogen  atoms  are  replaceable  by  metals;  we  should 
therefore  expect  to  find  four  classes  of  salts.  In  reality  only  two 
are  known,  M4P2O7  and  M2H2P2O7.  The  neutral,  as  well  as  the 
acid,  salts  of  the  alkalies  are  soluble  in  water;  the  neutral  salts  of 
other  bases  are  insoluble,  the  acid  salts  chiefly  soluble. 

Pyrophosphoric  acid  is  distinguished  from  the  ortho-acid  by 
the  fact  that  solutions  of  its  salts  give  a  white  precipitate,  Ag4P2O7, 
with  silver  nitrate,  and  from  the  meta-acid  by  not  coagulating 
albumen  and  giving  no  precipitate  with  barium  chloride. 

METAPHOSPHORIC  ACID,  HP03. 

148.  This  acid  is  obtained  by  heating  the  ortho-  or  the  pyro- 
acid  till  no  more  water  passes  off,  or  by  heating  ammonium  phos- 
phate  (NH4)2HP04.     Moreover,   on  dissolving  phosphorus  pent- 
oxide  in  cold  water,  the  product  is  at  first  chiefly  meta-acid. 

At  ordinary  temperatures  metaphosphoric  acid  is  a  vitreous 
solid  (hence  the  name  glacial  phosphoric  acid),  which  can  be  melted 
and  easily  drawn  out  into  threads.  On  being  heated  strongly 


218  INORGANIC  CHEMISTRY.  [§§  148- 

it  volatilizes  without  breaking  up  into  water  and  pentoxide.  When 
boiled  in  aqueous  solution  it  goes  over  into  orthophosphoric  acid. 
It  is  very  deliquescent;  use  is  made  of  this  property  occasionally. 

Metaphosphoric  acid  is  monobasic,  corresponding  to  the  formula 
HPOs.  Its  alkali  salts  only  are  soluble  in  water.  In  solution 
the  meta-acid  can  be  distinguished  from  the  ortho-  and  the  pyro- 
acids  by  its  ability  to  coagulate  albumen  and  give  white  precipi- 
tates with  chlorides  of  barium  or  calcium. 

The  vapor  of  this  substance  at  bright-red  heat  consists  chiefly 
of  H2P2Oe  molecules  (di-metaphosphoric  acid),  which  are  apparently 
liable  to  undergo  partial  dissociation  and  even  to  lose  a  small 
quantity  of  water. 

There  are  salts  of  various  acids  known,  which  must  be  regarded  as 
polymers  of  metaphosphoric  acid,  e.g.  K2P206,  potassium  di-metaphos- 
phate;  there  exist  also  tri-,  tetra-,  and  hexa-metaphosphates,  i.e.  salts 
of  the  acids  H3P3O9,  H4P4O12,  and  H6P6018. 

Hypophosphoric  Acid,  H4P2O6. 

149.  When  sticks  of  phosphorus  are  suspended  in  a  solution  of 
sodium  acetate  in  such  a  way  that  only  0.5  cm.  is  exposed  above 
the  level  of  the  liquid  and  the  temperature  is  kept  between  6° 
and  8°,  the  phosphorus  oxidizes  slowly  and  the  difficultly  soluble 
acid  sodium  salt  of  hypophosphoric  acid,  Na2H2P2O6 +6H2O  soon 
begins  to  crystallize  out.  It  can  be  purified  by  crystallization  from 
a  dilute  solution  of  acetic  acid.  If  this  salt  is  dissolved  in 
water  and  barium  chloride  added,  a  precipitate  of  barium  hypophos- 
phate  is  formed,  from  which  an  aqueous  solution  of  the  free  acid  can 
be  obtained  by  means  of  dilute  sulphuric  acid.  This  can  be  evaporated 
at  30°  to  a  sirupy  consistency  without  decomposition  and,  when  left 
in  a  vacuum,  yields  crystals  of  the  acid.  At  an  elevated  temperature 
and  in  the  presence  of  a  mineral  acid  phosphorous  and  phosphoric  acids 
are  formed.  This  behavior  justifies  the  consideration  of  hypophos- 
phoric acid  as  a  mixed  anhydride  of  the  two  last-named  acids: 

OH    HO      OH    OH 
OPOH    HOP  ->  OPOH   POH 
|OH__H]0     \o/ 

Phosphoric     Phosphorous     Hypophosphoric 
acid.  acid.  acid. 


151.]  PHOSPHOROUS  ACID.  219 

However,  it  has  not  yet  been  possible  to  prepare  the  hypophopphoric 
acid  by  melting  together  the  other  two  acids. 

From  the  determination  of  the  molecular  weight  of  the  methyl  ester 
it  seems  probable  that  the  formula  of  the  acid  is  H2PO3  and  not  H4P206 

Metaphosphorous  Acid,  HP02. 

150.  This  compound  was  discovered  by  VAN  DER  STADT  during  the 
slow  oxidation  of  phosphine  under  reduced  pressure  (§  137) : 

PH3  +  O2=H2  +  HPO2. 

The  sides  of  the  vessel  become  covered  with  feather-like  crystals  of  HP02. 
These  melt  at  a  much  higher  temperature  than  the  crystals  of  phos- 
phorous acid  and  are  converted  into  the  latter  by  the  action  of  water 
vapor. 

PHOSPHOROUS  ACID,  H3P03. 

151.  In  §  149  it  was  mentioned  that  this  acid  is  formed  by  the 
slow  oxidation  of  phosphorus  in  moist  air.     It  is  more  easily  pre- 
pared by  decomposing  phosphorus  trichloride  with  water: 

PC13  +  3H20  =  H3P03  +  3HC1. 

The  hydrochloric  acid  can  be  expelled  by  evaporating  at  180° 
and  the  phosphorous  acid  crystallizes  out  on  cooling. 

The  melting-point  of  phosphorous  acid  is  70.1°.  It  is  a  very 
hygroscopic  substance  Heating  decomposes  it  into  phosphoric 
acid  and  phosphine.  It  has  a  strong  reducing  action,  being  itself 
oxidized  to  phosphoric  acid.  The  oxygen  of  the  air  acts  on  it  very 
slowly.  It  precipitates  the  metals  from  solutions  of  gold  chloride, 
mercuric  chloride,  silver  nitrate,  etc.  A  characteristic  reaction  is 
the  reduction  of  sulphur  dioxide  to  sulphur,  which  takes  place  at 
ordinary  temperatures,  when  solutions  of  the  two  substances  are 
mixed. 

In  spite  of  its  three  hydrogen  atoms,  phosphorous  acid  acts  as 
a  dibasic  acid.  As  we  have  already  observed,  the  ionization  of 
poly  basic  acids  sometimes  affects  only  one  H*  ion  at  first,  the  others 
being  split  off  with  increasing  difficulty.  According  to  OSTWALD 
it  may  be  supposed  that  ionization  beyond  2H'  and  HPO3"  is  in 
this  case  so  difficult  that  the  acid  seems  to  be  only  dibasic.  The 
phosphites  are  not  oxidized  by  the  air,  but  they  yield  to  the  action  of 


220  INORGANIC   CHEMISTRY.  [§§152- 

oxidizing-agents;  e.g.  they  liberate  the  precious  metals  from 
their  salts,  as  does  also  the  acid  itself.  Heating  breaks  them  up 
into  hydrogen,  pyrophosphates  and  phosphide.  The  double 
phosphites  give  precipitates  with  baryta-  or  lime-water. 


Hypophosphorous  Acid,   H3P02. 

152.  Salts  of  this  acid  are  produced  by  heating  phosphorus  with  caustic 
soda,  lime-water  or  baryta-water  (§  136): 


3Ba(OH)2  +  8P  +  6H2O  =  3Ba(H2PO2)2  +  2 

It  can  be  set  free  from  these  salts  by  sulphuric  acid;  the  aqueous  solution 
is  concentrated  at  80°-90°  and  then  cooled  strongly,  whereupon  the  acid 
crystallizes  out.  Melting-point,  26.5°.  On  being  heated  at  130°-140° 
the  acid  splits  up  into  phosphorous  acid  and  phosphine;  at  a  somewhat 
higher  temperature  the  latter  acid  yields  phosphine  and  phosphoric  acid. 
The  equations  are: 

3H3PO2  =  2H3PO3  +  PH3  ;    3H3PO3  =  2H3PO4  +  PH3. 

Hypophosphorous  acid  is  a  very  strong  reducing-agent.  Gold,  silver 
and  mercury  are  precipitated  from  solutions  of  their  salts  by  the  free  acid 
as  well  as  its  salts.  Sulphur  dioxide  is  reduced  to  sulphur  at  ordinary  tem- 
peratures. In  these  reactions  the  acid  itself  is,  converted  into  phosphoric 
acid.  It  is  distinguished  from  phosphorous  acid  by  its  behavior  towards 
copper  sulphate  solution;  when  it  is  warmed  with  the  latter,  a  red  precipi- 
tate of  copper  hydride,  Cu2H2,  is  formed.  Hypophosphorous  acid  is  mono- 
basic. 

Compounds  of  Phosphorus  and  Sulphur. 

153.  Various  compounds  of  this  sort  are  known;  all  of  them  are 
obtained  by  warming  the  two  elements  together.  As  the  reaction  is 
very  vigorous  with  yellow  phosphorus,  the  red  form  is  usually  employed. 

The  compound  P2S5,  which  is  of  service  in  organic  chemistry,  is  a  yellow 
crystalline  substance,  melting  at  274°-276°  and  boiling  at  518°.  On  being 
warmed  with  water  it  yields  phosphoric  acid  and  sulphuretted  hydrogen. 
P2S5  unites  with  3  molecules  of  K2S  to  form  a  sulphophosphate,  K3PS4,  i.e.  a 
phosphate  whose  oxygen  is  replaced  by  sulphur. 

Several  compounds  containing  a  halogen  in  addition  to  phosphorus  and 
sulphur  are  known,  e.g.  PSC13.  This  phosphorus  sulphochloride  can  be  pre- 
pared by  treating  phosphorus  pentachloride  with  hydrogen  sulphide,  a  method 


155.]        COMPOUNDS  OF  PHOSPHORUS  AND  NITROGEN.       221 

analogous  to  that  of  forming  the  oxychloride  from  the  pentachloride  and 
water.  A  more  convenient  method  is  by  the  action  of  the  pentachloride 
on  the  pentasulphide,  which  carries  out  the  analogy,  to  oxy-compounds 
still  farther  (§  142)  : 

3PC15+P2S5=5PSC13. 

It  is  a  colorless  liquid,  boiling  at  125°.     Water  decomposes  it  into 
phosphoric  acid,  hydrochloric  acid  and  hydrogen  sulphide. 

Compounds  containing  Phosphorus  and  Nitrogen. 

154.  The  compounds  of  this  class  are  also  numerous.     Among  them 
are     amidophosphoric     acid,     ^Pi        >    and    diamidophosphoric    acid, 


OTT 
OP  />^jj  -,  .    As  their  names  indicate,  these  compounds  behave  like  acids. 

If  dry  ammonia  is  conducted  over  phosphorus  pentachloride,  a  white 
mass  is  obtained  which  consists  supposedly  of  ammonium  chloride, 
NH4C1,  and  a  compound  PC13(NH2)2.  With  water  it  forms  phosphamide, 
PO(NH)(NH2),  a  white  insoluble  powder.  On  being  boiled  with  water 
secondary  ammonium  phosphate  is  formed: 


PO(NH)(NH2) 

The  name  phospham  is  given  to  a  compound  P3H3N6,  which  is  formed 
from  the  product  of  the  action  of  ammonia  on  phosphorus  pentachloride, 
when  it  is  heated  in  the  absence  of  air  till  no  more  ammonium  chloride 
fumes  appear.  It  is  insoluble  in  water.  When  fused  with  potassium 
hydroxide,  it  breaks  up  as  follows: 

P3H3N6  +  9KOH  +  3H20  =  3K3P04  +  6NH3. 

By  the  interaction  of  P2S5  and  NH3  it  is  easy  to  obtain  a  com- 
pound P3N5,  phosphorus  nitride. 

ARSENIC. 

155.  Arsenic  occurs  in  nature  in  the  free  state  —  native.  More 
frequently  it  is  found  in  combination  with  sulphur  (realgar,  As2S2, 
and  orpiment,  As2Sc)  and  with  metals  (arsenopyrite,  or  mispickel, 
Fe  AsS,  and  cobaltiie,  Co  AsS)  ;  also  with  oxygen  as  As2O3  (arsenolite)  . 

The  extraction  of  the  element  from  these  minerals  is  simple 


222  INORGANIC  CHEMISTRY.  [§  155- 

Arsenopyrite  yields  arsenic  on  mere  heating,  the  latter  subliming. 
Arsenolite  is  reduced  with  carbon: 

2As203+6C  =  As  +  6CO. 

Physical  Properties  and  Allotropic  Conditions. — The  condition 
in  which  arsenic  usually  occurs  is  the  crystalline.  It  then  has  a 
steel-gray  color  and  a  specific  gravity  of  5.727  at  14°  and  is  a  good 
conductor  of  electricity.  It  sublimes  under  ordinary  pressure 
without  melting;  under  increased  pressure,  however,  it  melts  at 
500°.  By  sublimation  in  a  current  of  Hydrogen  a  second  crystal- 
lized form  can  be  obtained  together  with  a  black  modification, 
which  according  to  RETGERS  is  also  crystallized.  An  amorphous 
modification  results  from  the  decomposition  of  hydrogen  arsenide 
by  heat,  the  arsenic  appearing  as  a  dark  brown  deposit  on  the  sides 
of  the  glass.  Finally  there  is  a  yellow  modification  which  is  formed 
when  arsenic  vapor  is  condensed  in  a  dark  room  by  liquid  air. 
This  yellow  arsenic  is  very  sensitive  to  light;  even  at  the  tem- 
perature of  liquid  air  (— 180°) — at  which  it  is  stable  in  the  dark — 
it  is  converted  into  the  black  modification  by  the  light  of  a 
Welsbach  burner.  It  is  a  remarkable  fact  that  a  solution  of  the 
yellow  modification  in  chlorine  is  much  more  stable  toward 
light  and  heat  than  is  the  pure  substance.  Such  solutions  are 
obtained  in  concentrations  up  to  7%;  when  they  are  cooled 
yellow  arsenic  crystallizes  out.  The  relation  between  yellow  and 
black  arsenic  is  very  analogous  to  that  between  yellow  and 
red  phosphorus,  except  that  in  the  case  of  arsenic  the  yellow 
form  is  much  less  stable.  At  an  elevated  temperature  (360°) 
all  the  modifications  pass  over  into  the  ordinary  crystalline 
form. 

Vapor  Density. — The  lemon-yellow  vapor  of  arsenic  has  a 
density  of  10.2  (air=l)  at  about  860°,  which  makes  the  molecular 
weight  293.8.  At  1600°-1700°  the  vapor  density  is  less  by  half, 
being  5.40.  Since  the  atomic  weight  of  arsenic  is  75,  its  molecule 
therefore  contains  four  atoms  at  about  860°  and  two  at  1600°-1700°. 

Chemical  Properties. — Arsenic  is  not  affected  by  dry  air  at 
ordinary  temperatures;  in  moist  air  it  becomes  covered  with  a 
coating  of  oxide.  At  180°  it  burns  with  a  bluish  flame  to  the 
oxide  As4Oe,  giving  off  a  peculiar  garlic-like  odor.  At  an  elevated 


156.]  HYDROGEN  ARSENIDE.  223 

temperature  it  combines  with  many  elements  directly;    it  unites 
with  chlorine  without  the  aid  of  heat;  producing  scintillations. 

HYDROGEN  ARSENIDE.     ARSINE,  AsH3. 

156.  Direct  synthesis  from  the  elements  is  not  possible  with 
this  compound.  It  is  formed  when  almost  any  arsenic  compound 
comes  in  contact  with  nascent  hydrogen  (zinc  +  sulphuric  acid). 
When  thus  prepared  it  contains  considerable  hydrogen,  however. 
Pure  arsine  is  obtained  by  treating  zinc  arsenide  or  sodium  arsenide 
with  dilute  sulphuric  acid : 

As2Zn3 + 3H2S04= 2AsH3 + 3ZnSO4. 

Physical  Properties. — Hydrogen  arsenide  is  a  gas;  it  liquefies 
at  —40°,  but  does  not  solidify  as  low  as  —110°.  Sp.  g.  =  38.9 
(H=  1).  It  must  be  handled  with  great  care,  as  it  is  very  poison- 
ous. Fortunately  its  presence  can  be  easily  detected  by  its  peculiar, 
disagreeable  odor. 

Chemical  Properties. — Arsine  can  be  decomposed  into  its  ele- 
ments by  heat.  If  the  gas  is  passed  through  a  hot  glass  tube, 
arsenic  is  deposited  on  the  sides  in  the  form  of  a  metallic  mirror. 
Induction  sparks  also  decompose  it.  By  the  latter  means  it  can 
be  shown  that  the  resulting  volume  of  hydrogen  is  1J  times  as 
large  as  that  of  the  gas  itself,  in  accord  with  the  formula  AsH3. 
It  is  an  endothermic  compound, 

As  +  3H-AsH3  =  -36.7  Cal., 

and  has  been  made  to  explode  by  fulminating  mercury  (§  419). 
Hydrogen  arsenide  burns  with  a  pale  flame,  yielding  water  and 
arsenious  oxide,  As2O3,  if  sufficient  air  is  present;  if  such  is  not  the 
case,  or  if  the  flame  is  cooled,  arsenic  is  deposited.  On  heating 
potassium  or  sodium  in  the  gas,  an  arsenide,  AsK3  or  AsNa3,  is 
formed.  Hydrogen  arsenide  precipitates  the  yellow  compound 
AsAg3  •  3AgNO3  from  a  very  concentrated  solution  of  silver  nitrate : 

AsH3  +  6AgN03  =  As  Ag3  •  3  AgN03. 

This  is  decomposed  by  the  addition  of  water  into  arsenious  acid, 
nitric  acid  and  metallic  silver,  the  latter  being  deposited. 


224  INORGANIC  CHEMISTRY.  [§§  156- 

This  reaction  is  called  GUTZEIT'S  test.  It  is  usually  carried  out  in  the 
following  way:.  A  drop  of  50%  AgNO3  solution  is  placed  on  a  piece  of 
filter  paper  and  the  moist  spot  is  held  over  a  test-tube  containing  some 
zinc,  dilute  sulphuric  acid  and  the  substance  to  be  tested  for  arsenic. 
A  plug  of  cotton  is  inserted  near  the  top  to  protect  the  paper  from  being 
spattered  by  the  effervescing  solution.  If  arsenic  is  present,  the  spot 
becomes  yellow,  and  turns  black  when  moistened  with  water. 

Composition  of  Arsine. — If  arsine  is  passed  over  hot  copper 
oxide,  water  and  copper  arsenide  are  formed.  The  ratio  of 
hydrogen  to  arsenic  in  arsine  is  determined  from  this  reaction. 
For  1  part  (by  weight)  of  hydrogen  24.97  parts  of  arsenic  are 
obtained.  The  molecular  weight  of  the  compound,  as  found  from 
the  specific  gravity  (see  above),  is  77.9;  since  the  atomic  weight  of 
arsenic  is  75,  the  formula  of  arsine  must  be  AsH3. 

Detection  of  Arsenic. 

157.  The  majority  of  arsenic  compounds  are  very  poisonous.  Several 
of  them  are  of  practical  use  and  hence  are  on  the  market,  e.g.  white 
arsenic,  ^gA  (rat-poison);  orpiment,  As2S3;  Schweinfurt  green,  or 
copper  arsemle.  Poisonings  with  these  substances  happen  occasion- 
ally. Some  arsenic  compounds,  because  of  their  pretty  green  color, 
are  still  used,  though  much  less  than  formerly,  in  dyeing  tapestries, 
portieres,  and  the  like.  Rooms  in  which  these  are  hung  usually  con- 
tain particles  of  arsenical  matter,  which  are  injurious  to  the  health. 
Further,  a  certain  species  of  mould,  penicilliwn  hrevicaule,  which  is 
sometimes  found  in  such  tapestries,  has  the  power  of  generating  volatile 
and  very  poisonous  arsenic  compounds.  The  chemist  is  therefore  quite 
frequently  called  upon  to  analyze  a  given  sample  (of  dyed  materials 
or  the  like,  or  the  contents  of  a  stomach)  for  arsenic.  For  this  pur- 
pose a  method  has  been  devised  which  enables  him  to  detect  with  cer- 
tainty extremely  small  amounts  of  arsenic.  It  involves  the  following 
operations:  The  organic  substance  in  question  is  at  first  disintegrated 
as  well  as  possible,  usually  by  digestion  with  hydrochloric  acid  on  the 
water  bath,  a  little  potassium  chlorate  being  added  from  time  to  time. 
Thus  the  arsenic  compound  is  oxidized  to  arsenic  acid.  When  the 
chlorine  has  been  expelled  by  warming  and  the  liquid  has  been  filtered, 
hydrogen  sulphide  is  passed  in  for  some  time  at  a  temperature  of  about 
80°  to  precipitate  the  arsenic  as  sulphide.  The  sulphide  is  then  dis- 
solved in  nitric  acid  (in  case  the  presence  of  antimony  is  suspected  it 
must  first  be  removed);  this  solution  is  evaporated  to  dryness  to  get 
rid  of  the  excess  of  acid,  the  dry  residue  is  dissolved  in  water,  and  this 


157.]  DETECTION  OF  ARSENIC.  225 

liquid  is  then  tested  in  the  MARSH  apparatus,  a  simple  form  of  which 
is  shown  in  Fig.  38. 

This  consists  of  a  small  flask,  in  which  hydrogen  is  generated  from 
zinc  and  sulphuric  acid;  the  liquid  to  be  investigated  is  poured  down 
the  thistle-tube;  if  arsenic  is  present,  arsine  is  formed.  The  mixture 
of  hydrogen  and  arsine  is  dried  by  calcium  chloride  in  the  wide  tube 
and  then  enters  a  tube  of  hard  glass,  which  is  narrowed  at  several  places 
and  drawn  to  a  point  at  the  further  end.  As  the  gas  leaves  the  ta- 
pering end,  which  is  bent  upward,  it  is  lighted.  Thereupon  the  tube  is 
heated  with  a  flame  on  the  near  side  of  a  narrowed  place.  The  arsine 
is  broken  up  and  arsenic  is  deposited  as  a  bright  metallic  mirror  in  the 
narrowed  part.  From  the  extent  and  thickness  of  the  deposit  one  can 
estimate  the  number  of  milligrams  of  arsenic  present.  If  the  hydro- 
gen  arsenide  is  not  heated,  it  passes  on  to  the  flame  and  is  burned.  A 


FIG.  38.— MARSH  APPARATUS. 


cold  porcelain  dish  held  in  the  flame  is  soon  coated  with  a  deposit  of 
arsenic,  which  is  readily  soluble  in  sodium  hypochlorite  solution  (sodium 
arsenate  being  formed) .  This  solubility  enables  us  to  distinguish  arsenic 
from  antimony. 

Arsenic  is  very  widely  distributed,  although  in  small  amounts; 
hence  we  always  have  to  reckon  with  the  possiblity  of  traces  of  it  being 
present  in  the  reagents  and  glass  utensils  of  the  laboratory.  In  order 
to  test  this  a  "blank  experiment"  is  performed,  i.e.  all  the  operations 
are  carried  out  with  duplicate  amounts  of  the  required  chemicals  but 
without  the  addition  of  the  substance  to  be  analyzed.  Not  until  the 
materials  used  are  proved  to  be  free  from  arsenic  is  it  permissible  to 
use  them  in  an  actual  test. 

Whether  or  not  textile  fabrics  and  the  like  have  been  dyed  with 
Schweinfurt  green  (copper  arsenite)  can  be  determined  easily  by  the 
GUTZEIT  test.  Another  method  is  to  use  the  above-mentioned  peni- 
cillium  brevicaule.  This  is  cultivated  on  bread  which  is  soaked  with 


226  INORGANIC  CHEMISTRY.  L§§  157- 

the  liquid  to  be  tested  for  arsenic.  The  least  trace  of  the  latter  reveals 
itself  by  a  characteristic  garlic-like  odor,  caused  by  the  evolution  of 
arsenical  gases. 

Compounds  of  Arsenic  with  the  Halogens. 

158.  Three  arsenic-halogen  compounds  of  the  type  AsX5  are 
known;  viz.,  the  pentachloride,  AsCl5,  the  penta-iodide,  AsI5, 
and  the  pentafluoride,  AsF5.  Aside  from  these  only  compounds 
of  the  type  AsXs  are  known. 

Arsenic  trichloride,  AsCls,  can  be  obtained  by  direct  synthesis 
or  by  the  action  of  hydrochloric  acid  on  white  arsenic.  The  latter 
way  is  analogous  to  the  formation  of  metal  chlorides  from  the  oxide 
and  hydrochloric  acid.  This  compound  is  a  colorless  oily  liquid 
having  a  specific  gravity  of  2.205  (dS).  It  freezes  at  —18°  and 
boils  at  130.2°.  It  is  extremely  poisonous.  When  exposed  to  the 
air  it  throws  off  dense  white  fumes.  With  a  little  water  it  forms 
an  oxychloride,  As(OH)2Cl;  with  much  water  hydrochloric  acid 
and  arsenious  oxide.  In  this  latter  system  a  rise  of  temperature 
results  in  partial  re-formation  of  the  trichloride,  which  is  volatile 
with  the  water  vapor.  The  following  equilibrium  seems  to  exist: 

As2O3  +  6HCl  <=±  2AsCl3  +  3H2O. 

Oxygen  Compounds  of  Arsenic. 

Two  such  compounds  are  known:   As2O3,  arsenious  oxide,  and 
arsenic  oxide. 


ARSENIOUS  OXIDE,    As203. 

159.  Arsenious  anhydride  (commonly  called  "arsenic"  or 
"  white  arsenic")  is  found  in  nature.  It  is  formed  by  the  com- 
bustion of  arsenic  in  air  or  oxygen  and  by  the  oxidation  of  arsenic 
with  dilute  nitric  acid.  It  is  manufactured  commercially  by 
roasting  arsenical  ores;  the  oxide  volatilizes  and  is  condensed  in 
brick-walled  chambers,  where  it  collects  as  a  white  powder 
("arsenic  meal")  .It  is  refined  by  sublimation  from  iron  cylinders,, 

Physical  Properties.  —  Arsenious  oxide  is  an  odorless  solid,  that 
does  not  melt  under  ordinary  pressure,  but  sublimes.  Under 
higher  pressure  it  is  possible  to  melt  it.  At  800°  its  vapor  density 


159:]  ARSENIOUS  OXIDE.  227 

is  198  (O  =  16),  which  makes  the  molecular  formula  AS^G. 
Above  this  temperature  dissociation  begins  and  at  1800°  the  vapor 
density  corresponds  to  the  formula  As2O3.  By  the  ebullioscopic 
method  (elevation  of  the  boiling-point)  the  molecular  formula  has 
been  found  to  be  As^e  at  205°  also  (in  boiling  nitrobenzene). 

Various  Modifications. — Arsenious  oxide  is  known  in  a  vitreous 
form  as  well  as  in  crystals  of  the  regular  and  monoclinic  systems. 

The  vitreous  modification  is  produced  when  the  compound  is  sub- 
limed or  heated  to  the  sublimation-point.  Sp.  g.  =3.738.  After  stand- 
ing  for  some  time  at  ordinary  temperatures,  this  form  becomes  white 
like  porcelain  because  of  conversion  into  isometric  crystals.  The  latter 
form  is  better  obtained  by  dissolving  the  vitreous  modification  in  water 
or  hydrochloric  acid  and  letting  it  crystallize  out.  During  the  crystal- 
lization the  strange  phenomenon  of  bright  luminescence  is  observed, 
which  is  caused  by  the  breaking  of  the  crystals.  This  phenomenon, 
which  is  also  noticed  in  other  crystallizations,  is  called  tribolumines- 
cence.  The  transformation  of  the  amorphous  into  the  regular  variety 
is  accompanied  by  the  evolution  of  heat  (5.330  Cal.).  The  monoclinic 
form  is  obtained  by  conducting  the  crystallization  above  200°  instead 
of  at  ordinary  temperatures.  If  the  lower  half  of  a  sealed  glass  tube 
containing  arsenious  oxide  be  heated  above  400°,  it  will  be  found  after 
cooling  that  the  lower  heated  part  contains  vitreous,  the  middle  mono- 
clinic,  and  the  upper  octahedral,  arsenious  oxide. 

Since  the  transformation  of  amorphous  into  crystallized  arsenious  oxide 
takes  place  even  at  ordinary  temperatures  (rapidly  at  100°)  and  with 
the  evolution  of  heat,  the  octahedral  form  is  to  be  regarded  as  the  stable 
one  at  ordinary  temperatures;  the  glassy  form  is  only  able  to  exist 
at  these  temperatures,  because  the  velocity  of  transformation  is  then 
very  small.  According  to  the  above,  if  octahedral  arsenious  oxide  is 
gradually  warmed,  we  have  first  a  transformation  into  monoclinic  and 
then  another  into  amorphous  arsenious  oxide.  The  transition  tem- 
peratures have  not  yet  been  determined. 

Chemical  Properties. — Arsenious  oxide  is  easily  reduced  to 
arsenic ;  for  example,  by  heating  with  charcoal  or  nascent  hydrogen. 
It  is  also  easily  oxidized  to  arsenic  oxide  and  is  therefore  useful  as 
a  reducing-agent.  This  oxidation  can  be  brought  about  by  chlorine, 
bromine  (bromine- wa ter) ,  iodine  solution,  potassium  perman- 
ganate, strong  nitric  acid,  etc.  It  is  slightly  soluble  in  water;  the 
solution  has  a  salty  metallic  taste  and  a  weak  acid  reaction.  In 
acids  it  dissolves  much  more  easily,  because  it  acts  towards  them 
as  a  basic  oxide.  It  was  stated  above  (§  158)  that  a  solution  of 


228  INORGANIC  CHEMISTRY.  [§§  153- 

the  oxide  in  hydrochloric  acid  gives  off  arsenious  chloride.  White 
arsenic  is  a  rank  poison;  freshly  precipitated  ferric  hydroxide 
serves  as  an  antidote. 


ARSENIC    OXIDE, 

160.  This  compound  cannot  be  prepared  like  the  correspond- 
ing phosphorus  compound  by  burning  arsenic  in  the  air,  for  the 
oxidation  goes  no  farther  than  to  arsenious  oxide.  The  higher 
oxide  can  only  be  prepared  by  heating  arsenic  acid  in  the  air  : 

2H3AsO4-3H2O  =  As205. 

This  arsenic  anhydride  is  a  white  glassy  substance,  that  dis- 
solves in  water  slowly,  going  over  into  arsenic  acid.  By  heating 
with  carbon  it  is  easily  reduced  to  arsenic.  At  an  elevated  tem- 
perature it  breaks  up  into  oxygen  and  arsenious  oxide.  Its  molec- 
ular weight  is  not  known;  the  formula  As20s  is  simply  empirical. 


Oxyacids  of  Arsenic. 

Two  of  these  are  known:  arsenious  acid,  H3As03  (only  in 
aqueous  solution  and  salts)  and  arsenic  acid,  H3As04. 

ARSENIOUS  ACID,  H3As03. 

161.  This  acid  exists  in  the  aqueous  solution  of  the  anhydride. 
It  still  remains  to  be  discovered,  however,  which  hydrate,  H3AsO3, 
HAsC>2  or  some  other,  is  present.  On  evaporation  the  anhydride 
and  not  the  acid  separates  out.  This  acid  forms  three  classes  of 
salts,  according  as  one,  two,  or  three  of  its  hydrogen  atoms  are 
replaced  "by  metals;  it  is  therefore  tribasic.  Certain  salts  are 
known  which  are  derived  from  a  meta-arsenious  acid,  HAsO2- 

The  salts  of  the  alkalies  are  soluble  in  water;  those  of  the  other 
metals  are  not,  but  dissolve  easily  in  acids,  however.  A  neutral 
arsenite-  solution  gives  a  yellow  precipitate  of  silver  arsenite, 
Ag3As03,  with  silver  nitrate. 

The  solution  of  the  free  acid  is  easily  oxidized  to  arsenic  acid  by 
iodine  solution: 

H3AsO3  +  I2+H20   =    H3As04+2HL 

Such  a  solution  can  therefore  also  be  employed  for  the  titration  of  iodine 
(§  93). 


163.]  ARSENIC  ACID.  229 

ARSENIC  ACID,  H3As04. 

162.  This  acid  is  most  easily  obtained  by  the  oxidation  of  a 
solution  of  arsenious  acid  by  warming  it  with  nitric  acid.     On 
concentrating  the  solution  the  compound  2H3AsO4+H20  separates 
out  (below  15°) ;  this  substance  gives  off  its  water  of  crystallization 
at  100°  and  yields  orthoarsenic  acid,  H3AsO4,  which  crystallizes  in 
fine  needles.     When  heated  further  it  gives  off  water  (at  180°)  and 
goes  over  into  pyroarsenic  acid,  H^A^Oy,  which  separates  in  the 
form  of  hard  glistening  crystals.      On  being  heated  still  higher 
the  latter  compound  gives  up  another  molecule  of  water,   the 
final  product  being  white  crystalline  meta-arsenic  acid,  HAsO3. 
This  conduct  is  completely  analogous  to  that  of  phosphoric  acid; 
however,  metaphosphoric  acid  cannot  be  converted  into  the  anhy- 
dride by  heat  as  can  arsenic  acid  (§  160).    The  pyro-  and  meta- 
arsenic  acids  are  stable  only  in  the  solid  state;  when  treated  with 
water  they  are  converted  into  the  ortho  acid,  the  transformation 
being  much  quicker  than  with  the  corresponding  phosphorous  acids. 

Orthoarsenic  acid  is  easily  soluble  in  water.  Its  salts,  the 
arsenates,  exist  in  three  classes;  of  the  tertiary  only  those  of  the 
alkalies  are  soluble  in  water.  The  reactions  of  arsenic  acid  are 
very  similar  to  those  of  phosphoric  acid  (§  146) ;  in  this  case  also 
a  mixture  of  ammonia,  ammonium  chloride  and  magnesium  sul- 
phate (magnesia  mixture)  precipitates  a  white  crystalline  am- 
monium magnesium  salt,  Mg(NH4)AsO4+6H20.  Ammonium 
molybdate  produces  a  yellow  finely  crystalline  precipitate,  whose 
composition  and  appearance  correspond  to  those  of  the  phos- 
phorus compound.  The  precipitates  formed  with  silver  nitrate 
are,  however,  unlike  in  color:  Ag3P04  is  yellow,  Ag3AsO*  reddish 
brown. 

Sulphur  Compounds  of  Arsenic. 

163.  Three   are   known:    arsenic   disulphide   (realgar) , 
arsenic  trisulphide  (orpiment),  As2S3;  arsenic  pentasulphide, 

Arsenic  disulphide,  As2S2, 

occurs  in  nature  as  realgar  (§155).  It  forms  beautiful  ruby-red  crystals 
of  a  specific  gravity  of  3.5.  It  is  used  as  a  pigment.  It  is  manufactured 
artificially  by  fusing  sulphur  and  arsenic  together;  the  resulting  products 
vary  in  composition,  however. 


230  INORGANIC  CHEMISTRY.  [§§  163- 


ARSENIC  TRISULPHIDE, 

Arsenic  is  precipitated  from  the  acid  solution  of  arsenious 
oxide  by  sulphuretted  hydrogen  as  sulphide;  in  this  respect  too 
it  behaves  as  a  heavy  metal.  In  the  above  reaction  arsenic  tri- 
sulphide  is  deposited  as  an  amorphous  yellow  powder.  A  pure 
solution  of  arsenious  acid  gives  no  precipitate  with  sulphuretted 
hydrogen,  but  simply  a  yellow  liquid  (§  196).  Arsenic  trisulphide 
occurs  in  nature  as  orpiment  (§  155),  having  a  laminated  crystal- 
line structure;  it  owes  its  name  to  its  beautiful  golden  lustre. 
By  fusing  artificial  arsenic  trisulphide  a  product  is  obtained  which 
is  very  similar  to  the  natural  orpiment,  but  has  a  lower  specific 
gravity  (2.7  instead  of  3.4).  Commercially  the  trisulphide  is 
prepared  by  fusing  white  arsenic  with  sulphur;  the  product  still 
contains  the  oxide,  however,  and  is  therefore  poisonous.  Arsenic 
trisulphide  is  insoluble  in  water  and  in  acids. 


ARSENIC  PENTASULPHIDE, 

After  sulphuretted  hydrogen  has  been  led  into  a  warm  acidu- 
lated solution  of  arsenic  acid  for  some  time,  arsenic  is  precipi- 
tated as  an  amorphous  yellow  powder  of  the  composition  As2S5. 
The  latter  is  also  obtained  by  fusing  arsenic  trisulphide  with  the 
required  amount  of  sulphur.  In  the  absence  of  air  it  can  be 
sublimed  without  decomposition.  It  is  insoluble  in  water  and  in 
acids. 

SULPHO-SALTS  OF  ARSENIC. 

164.  The  trisulphide  and  the  pentasulphide  of  arsenic  dissolve 
easily  in  alkali  sulphides,  forming  salts  of  sulpho-acids: 

As2S3  +  3K2S  =  2K3AsS3  ; 

Pot.  sulph- 
arsenite. 

As2S5+3K2S=2K3AsS4. 

Pot.  sulph- 
arsenate. 

The  formation  of  these  sulpho-salts  can  be  regarded  as  analo- 
gous to  that  of  an  oxy-salt  from  a  basic  oxide  and  an  acid  anhv- 
dride,  e.g.  : 

BaO  +  SO3=BaSO4. 


ANTIMONY.  .  231 

The  trisulphide  and  the  pentasulphide  are  therefore  to  be  con- 
sidered as  sulpho-anhydrides  of  those  sulpho-acids. 

The   sulpharsenates   can  also  be   obtained   from  arsenic   tri- 
sulphide with  the  aid  of  an  alkali  polysulphide: 

As2S3+  K2S3=  2KAsS3. 

Pot.  sulpho- 
meta-arsenate. 

This  reaction  can  be  explained  by  supposing  that  the  arsenic 
trisulphide  is  converted  into  the  pentasulphide  by  the  excess 
of  sulphur,  just  as  the  trioxide  is  oxidized  to  the  pentoxide. 

They  are  also  produced  by  treating  an  arsenate  with  hydro- 
gen sulphide: 

=  K3AsS4+  4H20. 


The  sulpharsenates  and  sulpharsenites  of  the  alkalies  dis- 
solve easily  in  water  and  can  be  obtained  in  the  crystalline  form 
from  the  solution;  those  of  the  other  metals  are  insoluble.  The 
free  sulpho-acids  are  unknown.  On  the  addition  of  an  acid  to 
the  solution  of  a  sulpho-salt,  the  liberated  sulpho-acid  breaks  up 
into  hydrogen  sulphide  and  arsenic  tri-  or  pentasulphide. 

ANTIMONY. 

165.  Antimony  occurs  in  nature  in  stibnite,  Sb2S3,  as  well  as 
in  many  less  common  minerals.  Stibnite  was  known  to  the 
ancients.  In  Japan  it  is  found  in  magnificent  large  crystals. 
Antimony  was  frequently  employed  by  the  alchemists.  BASILIUS 
VALENTINUS  in  the  latter  part  of  the  fifteenth  century  described 
its  extraction  from  stibnite  in  a  monograph  entitled  "  The  tri- 
umphal car  of  Antimonium." 

The  element  is  at  present  obtained  from  stibnite  by  two 
processes.  In  one  the  mineral  is  roasted,  being  thus  transformed 
into  antimonious  oxide.  This  oxide  is  then  reduced  with  charcoal 
to  metallic  antimony: 

I.  2Sb2S3+902=2Sb2O3+6S02; 
II.  2Sb2O3+3C  =4Sb  +  3CO2. 

The  other  method  is  to  fuse  the  mineral  with  iron: 


232  INORGANIC  CHEMISTRY.  [§§  165- 

The  crude  antimony  thus  obtained  usually  still  contains  arsenic, 
lead,  sulphur,  etc.  It  can  be  refined  by  fusing  with  a  little  salt- 
petre, the  impurities  being  oxidized. 

Physical  Properties. — Antimony  is  silvery-white  and  has  a 
high  metallic  lustre  and  a  laminate-crystalline  structure  (rhombo- 
hedral) ;  as  a  result  of  the  latter  it  is  very  brittle  and  can  be  easily 
pulverized.  Sp.  g.  =  6.71-6.86.  Melting-point,  629°  •  boiling- 
point,  1440°.  MENSCHING  and  V.  MEYER  succeeded  in  determin- 
ing the  vapor  density  at  1437°,  i.e.  slightly  below  the  boiling-point, 
and  found  that  the  molecule,  unlike  that  of  phosphorus  or  arsenic, 
consists  of  less  than  four  atoms. 

Like  arsenic  antimony  has  a  black  and  a  yellow  modification;  the 
latter  is  obtained  by  passing  air  into  liquid  stibine,  cooled  to  —90°.  It 
is  even  less  stable  than  the  yellow  arsenic. 

Chemical  Properties. — At  ordinary  temperatures  the  element  is 
not  affected  by  the  air;  when  heated,  it  burns  with  a  bluish- white 
flame  to  the  trioxide.  It  combines  with  the  halogens  directly, 
producing  scintillations  (§  27).  It  is  dissolved  by  hydrochloric 
acid,  although  very  slowly,  with  the  evolution  of  hydrogen,  thus 
asserting  its  metallic  character.  Aqua  regia  dissolves  it  readily. 

Uses. — Antimony  is  a  constituent  of  various  alloys.  The  most 
important  of  these  is  type-metal,  from  which  printer's  type,  is 
made.  Its  approximate  compostion  is  lead  (50%),  antimony  (25%) 
and  tin  (25%). 

HYDROGEN  ANTIMONIDE,  STIBINE,  SbH3. 

166.  Stibine  is  formed  when  nascent  hydrogen  acts  on  a  solu- 
ble antimony  compound.  It  is  best  prepared  by  treating  an  alloy 
of  one  part  of  antimony  and  two  parts  of  magnesium  with  dilute 
hydrochloric  acid.  The  product  consists  principally  of  hydrogen, 
but  contains  10-14%  SbH3.  If  this  gas  mixture  is  passed  through 
a  U-tube  and  the  whole  is  plunged  in  liquid  air,  stibine  con- 
denses to  a  white  solid  mass,  that  soon  melts  after  the  tube  is 
removed  from  the  liquid  air.  It  vaporizes  to  a  relatively  stable 
gas.  The  least  trace  of  oxygen,  however,  causes  some  antimony 
to  be  deposited. 

If  an  electric  spark  is  passed  through  stibine  gas  it  explodes,, 
antimony  is  set  free  and  the  volume  of  hydrogen  liberated  is  found 
to  be  1J  times  that  of  the  stibine,  which  is  in  accord  with  the 
formula  SbH3.  It  is  also  decomposed  rapidly  by  heating  the 
containing  vessel  above  150°. 


167.]  HALOGEN  COMPOUNDS  OF  ANTIMONY.  233 

Stibine  has  a  characteristic  musty  odor,  quite  unlike  that  of 
phosphine  or  arsine. 

When  the  mixture  of  hydrogen  and  stibine  evolved  from  the 
alloy  of  antimony  is  heated,  as  in  the  MARSH  experiment  (§  157), 
it  produces  a  metallic  mirror  and,  when  ignited,  the  flame  gives  a 
spot  on  cold  porcelain  similar  to  that  of  arsenic,  but  differing  from 
the  latter  in  its  darker  color,  insolubility  in  hypochlorite  solution 
and  less  volatility  when  heated  in  a  current  of  hydrogen.  Stibine 
precipitates  a  black  powder  from  silver  solution,  consisting  of  a 
mixture  of  silver  and  silver  antimonide, 


The  decomposition  of  stibine  has  been  carefully  investigated  by 
STOCK.  He  arrived  at  the  conclusion  that  the  decomposition  velocity 
in  clean  glass  vessels  at  room  temperature  proceeds  at  first  with  extreme 
slowness  but  increases  more  and  more  as  more  antimony  separates  out. 
Furthermore,  mirrors  of  antimony  produced  by  heating  stibine  and 
mirrors  of  black  antimony  made  by  subliming  antimony  in  a  vacuum 
and  condensing  the  vapor  at  the  temperature  of  liquid  air,  and  mirrors 
of  sublimed  metallic  antimony,  all  had  different  effects.  The  effective- 
ness of  the  mirrors  varied  not  only  with  the  size,  but  in  large  measure  also 
with  the  form  of  the  antimony  surface. 

It  was  found  that  the  stibine  dissociation  in  the  layer  adsorbed  by 
antimony  was  proportional  to  the  mass  of  the  layer;  under  this  assump- 
tion the  progress  of  the  dissociation  could  be  calculated  theoretically, 
and  was  found  to  agree  well  with  the  experimental  results.  Since  the 
amounts  of  adsorbed  gas  depend  on  the  surface  tension  and,  therefore, 
on  the  form  of  the  adsorbing  surfaces,  the  explanation  of  the  influence 
of  the  different  antimony  mirrors  is  obvious. 

The  fact  that  the  walls  of  the  vessel  influence  the  velocity  of  a 
reaction  has  also  been  established  in  many  other  instances. 

Halogen  Compounds  of  Antimony. 

167.  Two  compounds  of  this  element  with  chlorine  are  known: 
SbCl3  and  SbCl5. 

Antimony  trichloride,  SbCls,  is  obtained  by  treating  antimony 
sulphide  or  oxide  with  concentrated  hydrochloric  acid.  It  forms 
a  colorless  laminar-crystalline  mass,  which  is  so  soft  that  it  was 
formerly  known  as  "antimony  butter"  (butyrwn  antimonii).  Its 
melting-point  is  73.5°  and  its  boiling-point  223.5°;  its  vapor 
density  7.8  (air  =  l)  makes  the  formula  SbCls. 

It    dissolves   in   water   containing   hydrochloric    acid.     Water 


234  INORGANIC    CHEMISTRY.  [§§167- 

decomposes  it,  forming  difficultly  soluble  oxy chlorides.  The 
composition  of  the  precipitate  depends  on  the  amount  and  the 
temperature  of  the  water  used  in  the  decomposition.  There  is 
evidence  of  the  existence  of  the  compounds  SbOCl  and 
Sb405Cl2  (=2SbOCl,  Sb2O3),  both  of  which  crystallize.  The  pre- 
cipitated oxychlorides  on  being  repeatedly  boiled  with  water  eventu- 
ally lose  all  their  chlorine  and  go  over  into  the  trioxide,  Sb2O3. 

Powder  of  Algaroth,  once  used  in  medicine,  is  obtained  by  the 
decomposition  of  antimony  trichloride  with  water  and  has  nearly 
the  same  formula  as  the  second  of  the  above-mentioned  oxy- 
chlorides. 

Antimony  pentachloride,  SbCl5;  is  prepared  by  heating  anti- 
mony in  a  current  of  chlorine  or  treating  fused  trichloride  with 
chlorine.  It  is  a  yellow,  fuming,  ill-smelling  liquid,  which  crys- 
tallizes at  —6°.  When  heated  it  dissociates  into  the  trichloride 
and  chlorine.  It  unites  with  water,  forming  SbCl5-H2O  and 
SbCl5«4H20.  Hot  water  decomposes  it  into  hydrochloric  and 
pyroantimonic  acids. 

Oxygen  Compounds  of  Antimony. 

168.  Three  are  known:  antimony  trioxide,  Sb2Os,  antimony 
tetr oxide,  Sb2O4,  and  antimony  pentqxide,  Sb2O5. 

Antimony  trioxide  occurs  as  a  mineral,  senarmontite.  It  can 
be  obtained  by  burning  antimony  in  the  air,  as  well  as  by  the 
oxidation  of  antimony  with  dilute  nitric  acid.  It  is  dimorphic, 
occurring  in  both  regular  and  rhombic  crystals. 

It  is'x^a  light  yellow  crystalline  powder,  almost  insoluble  in 
water.  It  volatilizes  at  1560°;  the  vapor  density  at  this  tem- 
perature corresponds  to  the  formula  Sb^e-  It  is  insoluble  in 
sulphuric  and  nitric  acids  but  easily  soluble  in  hydrochloric  and 
tartaric  acids  and  in  alkalies.  On  being  heated  in  the  air  it  turns 
to  the  tetroxide. 

The  corresponding  hydroxide  is  Sb(OH)3.  This  hydrate  sepa- 
rates out  when  tartar  emetic  (see  below)  is  decomposed  with 
dilute  sulphuric  acid.  It  gives  up  one  molecule  of  water  readily 
and  passes  over  into  the  hydroxide  SbO -OH,  meta-antimonious  acid. 

The  latter  is  more  easily  obtained  by  treating  a  solution  of  the 
trichloride  with  soda  solution: 


169.]        ANTIMONY  PENTOXIDE    AND  ANTIMONIC  ACID.       235 
2SbCl3 + 3Na2CO3 + H2O = 2SbO  •  OH + GNaCl + 3CO2. 

It  appears  as  a  white  precipitate,  which  is  converted  into  an- 
timonic  oxide  by  boiling  with  water.  This  meta-antimonious 
acid  is  dissolved  by  alkalies,  forming  salts  of  the  acid.  One  of 
them  which  has  been  obtained  crystallized  is  the  sodium  meta- 
antimonite,  NaSbO2+3H2O.  The  latter  is  difficultly  soluble  in 
water,  and  decomposes  on  concentration  of  its  solution. 

On  the  other  hand,  antimony  hydroxide  displays  basic  proper- 
ties by  uniting  with  acids  to  form  salts.  There  are  salts  known  of 
Sb(OH)3,  as  well  as  of  SbO -OH.  Examples  of  the  former  kind* 
are  the  crystallized  antimony  sulphate,  Sb2(SO4)3,  and  the  nitrate, 
Sb(NO3)3.  In  analogy  with  other  trivalent  metals  double  salts 
are  known,  e.g.  KSb(SO4)2.  As  to  the  salts  derived  from  SbO-OH, 
we  may  look  upon  the  group  SbO  as  taking  the  place  of  a  uni- 
valent  metal.  Thus  SbO-OH  may  be  compared  with  KOH.  For 
this  reason  the  group  (SbO)  has  been  given  the  name  antimonyl; 
one  of  its  salts  is  antimonyl  sulphate,  (SbO)2S04.  The  most 
familiar  antimonyl  compound  is  tartar  emetic,  potassium  antimonyl 
tartrate, 

(SbO)  ^4^-4^6  ~^~  i^2O, 

which  is  employed  in  medicine.     See  ORG.  CHEM.  §  192. 
ANTIMONY  PENTOXIDE  AND  ANTIMONIC  ACID. 

Antimonic  acid,  H3Sb04,  is  obtained  by  warming  antimony 
with  concentrated  nitric  acid  and  also  by  decomposing  the  penta- 
chloride  with  water.  It  is  a  white  powder,  almost  insoluble  in 
water  and  nitric  acid;  nevertheless,  when  moist,  it  turns  litmus 
paper  red.  On  heating  saltpetre  with  powdered  antimony  the 
potassium  salt  of  meta-antimonic  acid,  KSb03,  is  formed  in  an 
explosive  reaction.  When  this  is  boiled  with  water  it  dissolves, 
producing  monopotassium  orthoantimoniate,  KH2SbO4;  on  fusing 
with  potash  potassium  pyroantimoniate,  K4Sb2Oy,  is  formed,  which 
dissolves  in  water,  giving  2KOH  and  K2H2Sb2O7+6H20.  In  the 
case  of  antimony,  as  in  that  of  phosphorus,  we  meet  with  three  kinds 
of  acids  belonging  to  the  highest  stage  of  oxidation:  their  formulae 
correspond  to  those  of  the  analogous  phosphorus  compounds. 


236  INORGANIC  CHEMISTRY. 

Antimony  pentoxide,  Sb20s  (molecular  weight  unknown),  can 
be  obtained  by  heating  antimonic  acid  at  300°.  It  is  a  yellow 
amorphous  powder,  soluble  in  hydrochloric  acid.  If  heated 
strongly  it  gives  up  part  of  its  oxygen  and  goes  over  into  anti- 
mony tetroxide,  Sb2C>4,  a  white  powder  that  turns  yellow  on 
heating  but  resumes  its  original  color  on  cooling.  This  tetroxide 
can  be  regarded  as  antimonyl  meta-antimoniate,  SbOa-SbO. 

Sulphur  Compounds  of  Antimony. 

169.  Antimony  trisulphide,  Sb2S3,  is  found  in  nature  (§  165). 
It  can  be  made  by  leading  hydrogen  sulphide  into  a  hydrochloric 
acid  solution  of  the  trichloride,  from  which  it  is  deposited  as  an 
amorphous  red  powder.     It  can  be  melteci;  on  cooling  it  crystal- 
lizes and  takes  on  the  appearance  of  stibnite. 

Antimony  pentasulphide,  Sb2S5,  is  precipitated  when  hydrogen 
sulphide  is  passed  into  the  acidified  solution  of  antimonic  acid. 
It  is  more  easily  obtained  by  the  decomposition  of  sodium  sulph- 
antimoniate  with  dilute  sulphuric  acid.  It  forms  an  amorphous 
orange-red  powder,  which  splits  up  into  sulphur  and  the  trisulphide 
on  being  strongly  heated.  It  is  insoluble  in  dilute  acids;  boiling- 
iiot  concentrated  hydrochloric  acid  dissolves  it,  forming  antimony 
trichloride,  hydrogen  sulphide  and  sulphur.  In  aqueous  solutions 
of  alkalies  and  their  sulphides  it  dissolves  easily  with  the  formation 
of  sulphantimoniates,  M3SbS4  The  best  known  of  these  is  sodium 
sulphantimoniate,  Na3SbS4+9H20  ("  SCHLIPPE'S  salt").  It  can 
be  obtained  by  boiling  antimony  trisulphide  with  sulphur  and 
caustic  soda  solution.  It  crystallizes  in  large  colorless  tetrahedrons, 
is  easily  soluble  in  water  (1  part  by  weight  in  2.9  parts  water  at 
15°)  and  reacts  alkaline.  It  is  decomposed  by  acids,  depositing 
pentasulphide;  even  carbonic  acid  causes  this,  hence  the  crystals 
become  covered  with  a  yellowish-red  coating  of  pentasulphide  after 
having  stood  some  time  in  the  air.  The  free  sulphantimonic  acid 
is  not  known. 

BISMUTH. 

170.  This  elejnent  belongs  undoubtedly  among  the  metals,  so 
far  as  its  physical  character  is  concerned;   its  chemical  properties 
also  class  it  with  them  in  almost  every  respect,  inasmuch  as  its 
oxides  are  mainly  basic  in  their  behavior. 


171.]  COMPOUNDS  OF   BISMUTH.  237 

It  is  found  chiefly  in  the  native  state;  but  a  sulphide,  Bi2Ss, 
bismuth  glance  and  a  telluride,  tetradymite,  also  occur  in  nature. 
Bismuth  is  obtained  from  the  latter  by  roasting  to  the  oxide  Bi203 
and  reducing  with  charcoal.  The  native  metal  is  usually  very 
pure.  If  refining  is  necessary,  the  fused  metal  is  allowed  to  flow 
over  a  hot,  somewhat  inclined  iron  plate,  so  that  the  impurities 
are  oxidized.  The  amount  of  bismuth  found  in  nature  is  not 
very  great. 

Physical  Properties. — Bismuth  is  externally  very  similar  to 
antimony;  it  is  crystallized  and  very  brittle  and  has  a  metallic 
lustre,  but  differs  from  antimony  in  having  a  reddish-white  color. 
Sp.  g.  =  9.823.  It  melts  at  286.3°  and  boils  at  1420°.  it  can 
be  distilled  in  a  current  of  hydrogen. 

Chemical  Properties. — At  ordinary  temperatures  bismuth  is 
unaffected  by  the  air.  On  being  heated  it  turns  to  the  trioxide. 
It  combines  with  the  halogens  directly.  It  is  not  attacked  by 
hydrochloric  or  sulphuric  acid  at  ordinary  temperatures,  but  nitric 
acid  dissolves  it  readily  to  form  the  nitrate.  On  being  heated  with 
sulphuric  acid,  it  gives  off  sulphur  dioxide  and  forms  the  sul- 
phate. No  hydrogen  compound  of  bismuth  is  known. 

Bismuth  is  employed  in  the  manufacture  of  easily  fusible  alloys 
such  as  are  used  in  making  casts  of  woodcuts,  stereotypes,  etc. 
The  most  common  of  these  alloys  are  NEWTON'S  metal  (8  bismuth, 
5  lead,  3  tin;  melting-point  94.5°),  ROSE'S  metal  (2  bismuth, 

1  lead,  1  tin;  melting-point  93  75°)  and  WOOD'S  metal  (4  bismuth, 

2  lead,  1  tin,  1  cadmium;    melting-point  60.5°). 

Halogen  Compounds. 

171.  Compounds  of  the  type  BiX3  only  are  known.  Bismuth 
chloride,  BiCl3,  is  formed  by  direct  synthesis  from  the  elements, 
but  more  easily  by  dissolving  bismuth  in  aqua  regia.  It  is  white 
and  crystallized.  Its  melting-point  is  between  225°  and  230°  and 
its  boiling-point  at  435°.  Its  vapor  density,  11.35  (air=l),  gives 
it  the  formula  BiCl3.  On  being  dissolved  in  a  little  water  it  forms  a 
sirupy  liquid;  an  excess  of  water  gives  bismuth  oxychloride,  BiOCl, 
and  hydrochloric  acid.  This  oxychloride  is  a  white  powder,  in- 
soluble in  water  but  soluble  in  acids. 


238  INORGANIC  CHEMISTRY.  [§§  172, 

Oxygen    Compounds. 

172.  Three  oxides  are  known:   BiO,  Bi2O3,  and  Bi02. 

Bismuthous  oxide,  BiO,  is  obtained  by  adding  an  alkaline  stannous 
chloride  solution  to  a  solution  of  bismuth  chloride.  It  is  deposited 
as  a  dark-brown  precipitate  of  BiO.  When  heated  in  the  air  it  smolders 
like  tinder.  It  is  doubtful  whether  this  precipitate  is  a  homogeneous 
substance  or  a  mixture  of  Bi203  with  finely  divided  bismuth. 

Bismuth  trioxide,  Bi2O3,  is  the  most  familiar  oxide  of  this  ele- 
ment. It  has  strictly  basic  properties.  In  order  to  prepare  it  we 
can  heat  the  nitrate  or  carbonate  or  we  can  precipitate  the  hydroxide 
from  the  solution  of  a  bismuth  salt  by  means  of  a  base  and  heat 
the  precipitate.  If  a  boiling  solution  of  a  bismuth  salt  is  treated 
with  caustic  potash,  the  trioxide  separates  out  in  glistening  needles 
of  microscopic  dimensions.  Like  the  corresponding  oxides  of 
arsenic  and  antimony,  it  is  dimorphic. 

Bismuth  dioxide,  BiO2,  has  been  little  studied;  it  is  a  reddish-yellow 
powder. 

Bismuth  pentoxide,  Bi205,  appears  as  a  reddish-brown  powder,  which 
is  very  unstable  and  evolves  oxygen  on  heating,  as  it  also  does  when 
warmed  with  sulphuric  acid.  Hydrochloric  acid  does  not  convert  it 
into  the  corresponding  pentachloride,  BiCl5,  but  produces  the  trichloride 
BiCl3  and  free  chlorine. 

Hydroxides  and  Salts. 

173.  Bismuth  hydroxide,  Bi(OH)3,  is  obtained  by  precipitating 
a  bismuth  salt  with  an  alkali.  It  is  an  amorphous  white  powder, 
insoluble  in  potassium  hydroxide  or  ammonia.  At  100°  it  goes 
over  into  the  compound  BiO -OH  with  the  loss  of  a  molecule  of 
water.  Both  of  these  hydroxides  are  wholly  basic  in  character. 
The  salts  derived  from  Bi(OH)3  are  called  neutral,  those  from 
BiO -OH  basic. 

The  neutral  nitrate,  Bi(N03)3,  is  obtained  by  dissolving  bis- 
muth in  nitric  acid.  It  crystallizes  with  five  molecules  of  water 
in  large  translucent  triclinic  prisms.  It  is  deliquescent.  The 
addition  of  much  water  converts  it  into  the  basic  nitrates,  several 
of  which  are  known.  By  treating  the  neutral  nitrate  with  about 
20  parts  of  boiling  water  a  product  is  obtained  whose  composition 
is  not  perfectly  constant  for  different  preparations,  but  corresponds 
nearly  to  the  formula  (BijAOe-^OsVC^O^,  or  2BiONO3 


175.] 


SUMMARY  OF    THE  NITROGEN  GROUP. 


239 


+  Bi(NO3)3  +  3Bi(OH)3.     This  is  the  bismuth  subnitrate,  which 
is  used  in   medicine. 

Bismuth  sulphate,  Bi2(S04)3,  is  obtained  as  an  amorphous 
white  substance  when  the  metal  is  heated  with  concentrated  sul- 
phuric acid.  With  water  it  forms  a  basic  sulphate,  Bi2(OH)4SO4. 

Sulphur  Compounds. 

174.  Bismuth  trisulphide  is  found  in  nature  (§  170);  artificially 
it  can  be  prepared  by  heating  bismuth  with  sulphur  or  by  leading 
hydrogen  sulphide  into  the  aqueous  solution  of  a  bismuth  salt. 
In  the  latter  case  it  comes  down  as  an  amorphous  black  powder 
that  is  easily  soluble  in  warm  dilute  nitric  acid.  It  is  insoluble 
in  alkalies  and  their  sulphides,  hence  forms  no  sulpho-salts.  When 
heated  with  an  alkali  sulphide  solution  to  200°  it  takes  on  a  crys- 
talline form  similar  to  that  of  the  natural  mineral. 


SUMMARY  OF  THE  NITROGEN  GROUP. 

175.  Like  the  halogens  and  the  elements  of  the  oxygen  group, 
the  elements  just  discussed,  viz.  nitrogen,  phosphorus,  arsenic, 
antimony  and  bismuth,  also  form  a  natural  group.  Their  family 
relation  shows  itself  even  in  the  formula  types  of  their  compounds. 
The  hydrogen  compounds  have  the  type  RH3  (lacking  with  bis- 
muth), the  halogen  compounds  RX3  and  RX5  (the  latter  also  lack- 
ing with  bismuth),  the  oxygen  compounds  R2Os  and  R20s.  In 
other  words,  the  elements  of  this  group  are  trivalent  or  pentivalent. 
We  find  here,  just  as  in  the  groups  previously  studied,  that,  as  the 
atomic  weight  increases,  a  gradual  change  occurs  in  the  physical 
properties.  This  is  shown  by  the  following  small  table: 


N. 

P. 

As. 

Sb. 

Bi. 

Atomic  weight.  .  . 
Specific  gravity.  .  . 
(  Water  =  1) 
Melting-point.  .  .  . 
Boiling-point  
Color  

14.01 
0.885 
liquid 

-194.4° 
colorless 

31.0 
1.8-2.1 

+  44.4° 
+  278° 
yellow  or  red 

74.96 
4.7-5.7 

ca.  800° 
gray 

120.2 
6.7 

629° 
1440° 
white 

208.0 

9.8 

286° 
1420° 
pink 

In  the  chemical  properties,  also,  regular  variations  are  to  be 
observed,  all  of  which  can  be  summed  up  in  the  general  statement 


240  IXORGANIC  CHEMISTRY.  [§§175- 

that  the  metalloid  character  gives  way  to  the  metallic  character 
as  the  atomic  weight  increases.  Nitrogen  forms  either  indifferent 
or  acid-forming  oxides  only;  so  does  phosphorus;  arsenic,  on  the 
contrary,  displays  a  very  feebly  basic  character  in  arsenious  oxide, 
since  this  oxide  forms  the  trichloride  with  hydrochloric  acid,  the 
trichloride  reacting  inversely  with  water,  however,  and  breaking 
up  into  hydrochloric  acid  and  arsenious  oxide.  In  antimony 
trioxide  this  basic  character  is  a  little  stronger;  some  salts  and 
double  salts  of  it  with  acids  are  known.  The  corresponding  chlo- 
ride does  not  suffer  an  immediate  hydrolytic  dissociation  with 
water,  but  oxychlorides  are  formed,  which  require  a  great  deal 
of  water  to  convert  them  entirely  into  the  trioxide.  While  the 
highest  oxides  of  arsenic  and  antimony  have  strictly  acid  prop- 
erties, with  bismuth  the  acidic  nature  has  practically  disappeared; 
the  oxide  Bi2Os  has  exclusively  basic  properties  and  the  higher 
oxide  Bi2O5  acts  like  a  peroxide,  giving  off  oxygen  readily  (it 
generates  chlorine  with  hydrochloric  acid)  and  going  over  into 
the  lower  oxide  Bi2O3.  Bismuth  trichloride,  Bids,  gives  the  oxy- 
chloride,  BiOCl,  with  water  and  this  is  not  decomposed  by  an 
excess  of  water. 

In  the  hydrogen  compounds,  too,  the  gradual  change  of  the 
properties  is  very  apparent.  Consider  the  stability  for  example: 
ammonia  requires  a  very  high  temperature  for  decomposition; 
phosphine  and  arsine  a  much  lower  temperature;  stibine  is  unstable 
at  ordinary  temperatures  when  it  comes  in  contact  with  oxygen, 
and  the  hydrogen  compound  of  bismuth  is  so  unstable  that  the 
conditions  for  its  formation  and  existence  have  not  yet  been 
ascertainable.  A  similar  change  is  noticeable  in  their  ability 
to  form  XELj  ions  in  aqueous  solutions;  it  is  strong  in  ammonia, 
much  weaker  in  phosphine  and  wholly  absent  in  arsine  and 
stibine. 

In  the  sulphur  compounds  a  progressive  change  of  color  is 
observed.  P2S5  is  bright  yellow,  As2S5  deep  yellow,  Sb2$5  red 
and  Bi2Ss  black.  The  first  three  are  sulpho-anhydrides  of  suipho- 
acids  (§  164);  bismuth  sulphide  is  not,  however,  thus  displaying 
again  the  more  basic  nature  of  bismuth. 


176.J  ALLOTROPIC  FORMS  OF  CARBON.  241 


CARBON. 

176.  Carbon  occurs  in  nature  both  free  and  combined.  In 
combination  it  is  found  in  large  quantities  in  the  salts  of  carbonic 
acid,  above  all  in  calcium  carbonate,  limestone,  which  is  of  the 
widest  occurrence  and  is  even  known  to  form  great  mountains. 
Farther,  carbon  is  one  of  the  constituent  elements  of  animals  and 
plants.  It  is  found  in  these  in  numerous  compounds.  Still  larger 
is  the  number  of  artificially  prepared  carbon  compounds.  The  com- 
pounds of  carbon  exceed  in  number  all  other  compounds  together. 
For  this  reason  and  because  of  the  peculiarities  of  the  carbon 
compounds  it  is  customary  to  treat  them  by  themselves,  as  "  organic 
chemistry."  However,  that  we  may  be  able  to  obtain  a  general 
survey  of  the  elements,  it  is  deemed  advisable  to  discuss  certain 
compounds  of  carbon  in  inorganic  chemistry  as  well. 

Allotropic  Forms   of  Carbon. 

We  know  of  three:   diamond,  graphite  and  amorphous  carbon. 

(a)  Diamond. — LAVOISIER  found,  in  1773,  that  this  mineral 
can  be  burned  to  carbon  dioxide.  In  1814  DAVY  proved  that, 
when  diamond  burns,  nothing  else  than  this  gas  is  formed,  so 
that  diamond  must  be  pure  carbon.  Furthermore  when  the 
carbon  dioxide  given  off  by  the  combustion  of  diamond  is  absorbed 
by  sodium  hydroxide,  a  soda  is  produced  which  is  in  every  respect 
identical  with  ordinary  soda.  Indeed,  it  has  been  found  possible 
to  manufacture  diamonds  from  amorphous  carbon  (see  below). 

The  diamond  crystallizes  in  the  isometric  system.  Usually  it 
is  colorless,  but  yellow  and  black  diamonds  are  also  known;  the 
black  ones  are  called  carbonado.  The  specific  gravity  of  diamond 
is  3.50-3.55.  It  is  a  poor  conductor  of  heat  and  electricity.  The 
refractive  index  is  very  high:  n=2A2.  The,  diamond  is  so  hard 
that  it  scratches  all  other  substances.  If  it  is  subjected  to  a 
very  high  temperature  in  the  absence  of  air,  it  gradually  turns  to 
graphite.  It  resists  the  action  of  the  strongest  oxidizing-agents, 
e.g.  a  mixture  of  nitric  acid  and  potassium  chlorate. 

In  1893  MOISSAN  succeeded  in  making  diamonds  artificially, 
although  they  were  very  small,  the  largest  being  about  0.5  mm.  in 


242  INORGANIC  CHEMISTRY.  [§176- 

> 
diameter.     His  method  consists  essentially  in  dissolving  carbon 

in  molten  iron  at  a  high  temperature  and  then  cooling  it  rapidly. 


FIG.  39. — ARTIFICIAL  DIAMONDS  (MAGNIFIED). 

This  is  accomplished  as  follows:  Iron  is  brought  in  contact  with 
pure  carbon  (sugar  charcoal)  in  the  electric  furnace  at  a  high  tem- 
perature. After  the  iron  has  become  saturated  with  carbon  at 
about  3000°,  the  fused  mass  is  suddenly  cooled  by  pouring  it  into 
a  hole  drilled  in  a  copper  block,  which  is  kept  cold  by  water,  and 
at  once  covering,  the  cavity  with  an  iron  stopper.  When  the  iron 
is  all  cold  it  is  dissolved  away  by  acids,  leaving  the  carbon  which 
did  not  combine  with  the  iron.  This  residual  carbon  consists 
partly  of  small  diamonds,  which  are  identical  with  the  natural 
diamond  in  hardness,  crystal  form,  etc.  Fig.  39  presents  an 
enlarged  view  of  some  artificial  specimens;  they  display  the  same 
properties  as  the  rough  natural  diamonds,  particularly  the  rounded 
edges  and  angles  and  the  stria tions. 

The  formation  of  the  diamond  by  this  method  has  been  explained 
by  BAKHUIS  ROOZEBOOM  as  follows :  In  all  probability  the  transition 
of  diamond  into  graphite  is  endothermic.  For  this  reason  diamond 
is  the  more  stable  form  at  lower  temperatures,  graphite  at  higher 
ones,  in  analogy  to  the  rhombic  and  monoclinic  modifications  of 
sulphur.  But,  while%the  velocity  of  transformation  of  monoclinic 
sulphur  is  fairly  great  at  low  temperatures  and  the  monoclinic 
sulphur  can  thus  exist  only  for  a  short  time  below  its  transition 
point,  the  transition  velocity  of  graphite  into  diamond  is  practically 
zero  for  temperatures  below  1000°.  Carbon  that  has  crystallized 
from  molten  iron  in  the  form  of  graphite  cannot,  therefore,  pass  over 
into  diamond.  The  rapid  cooling  of  the  molten  iron,  however, 


176.] 


ALLOTROPIC  FORMS  OF  CARBON. 


243 


has  the  effect  of  bringing  the  carbon  into  the  region  of  temperature 
in  which  diamond  is  the  stable  modification;  it  can  therefore 
separate  in  this  form  from  its  solution. 

The  electric  furnace  that  MOISSA.N  used  for  these  and  numerous 
other  experiments  is  very  simple  in  construction.  It  consists  of  two 
blocks  of  unslaked  lime  that  fit  tightly  together.  In  the  lower  block 


FIG.  40. — MOISSAN'S  ELECTRIC  FURNACE  (CROSS-SECTION). 

there  is  a  trough  in  which  the  carbon  terminals  are  laid.  The  upper 
block  is  slightly  hollowed  out  on  its  lower  side  so  as  to  reflect  the  heat 
rays  on  to  the  crucible.  Fig.  40  shows  a  cross-section  of  an  electric 
furnace,  Fig.  41  a  picture  of  the  same  apparatus  in  operation. 

The  temperatures  obtained  in  the  electric  furnace  are  as  follows: 

Current  of   30  amperes  and  55  volts  with  a  steam-engine  of   4  H.P.,  2250° 


100 
450 


45 
70 


50 


2500° 
3000° 


FIG.  41. — MOISSAN'S  FURNACE  IN  OPERATION.     (AFTER  MOISSAN.) 


244  INORGANIC  CHEMISTRY.  [§§  176- 

The  last-named  temperature  can  however  only  be  maintained  for  .a 
brief  period,  as  the  unslaked  lime  soon  melts  and  flows  like  water.  At 
2500°  the  lime  becomes  crystalline  in  structure  after  a  few  minutes. 

(6)  Graphite  is  also  crystallized  carbon.  Unlike  diamond,  it  is 
very  soft  and  opaque  and  a  good  conductor  of  heat  and  electricity. 
Sp.  g.  =2.09-2.23.  As  was  stated  above,  graphite  can  be  pre- 
pared artificially  by  the  crystallization  of  carbon  from  molten  iron 
and  by  heating  diamond  strongly.  There  are  various  kinds  of 
graphite.  If  graphite  is  treated  with  a  mixture  of  perfectly  dry 
potassium  chlorate  and  very  concentrated  nitric  acid,  it  turns  to  a 
yellow  crystallized  substance  containing  hydrogen  and  oxygen,  in 
addition  to  carbon,  and  called  graphitic  acid.  This  substance  is 
peculiar  in  that  it  decomposes  explosively  on  heating  and  yields  a 
large  volume  of  extremely  fine  amorphous  carbon.  Graphite  is 
used  in  the  manufacture  of  lead  pencils,  crucibles,  electrodes, 
polishes,  etc. 

(c)  Amorphous  Carbon. — This  is  obtained  !n  the  purest  state 
by  charring  sugar.  The  resulting  mass  is  boiled  with  acids  to 
remove  the  mineral  matter  and  finally  heated  red-hot  in  a  current 
of  chlorine  for  quite  a  while  to  remove  all  the  hydrogen.  It  can 
also  be  prepared  from  soot.  Amorphous  carbon  is  opaque,  black 
and  infusible.  At  the  highest  temperature  that  MOISSAN  could 
reach  with  his  furnace  by  employing  a  current  of  2000  amperes 
and  80  volts  (obtained  with  a  300  horse-power  engine)  it  was  barely 
possible  to  make  carbon  sublime.  The  sublimate  was  graphite. 
Amorphous  carbon  has  a  specific  gravity  of  1.5-2.3. 

Various  sorts  of  amorphous  carbon  are  known.  They  are 
probably  different  allotropic  modifications,  or  mixtures  of  such. 
Gas  carbon  and  coke  are  obtained  as  residues  in  the  dry  distilla- 
tion of  coal.  They  conduct  heat  and  electricity.  Wood  charcoal 
is  very  porous  and  can  condense  large  quantities  of  gases  in  its 
pores,  e.g.  90  times  its  own  volume  of  ammonia  (see  also  §  111). 
When  warmed  or  when  the  pressure  is  reduced,  these  gases  all 
escape  again.  Bone-black  is  obtained  by  heating  bones  away 
from  air;  the  resulting  black  mass  is  treated  with  hydrochloric 
acid  to  remove  the  phosphates  and  carbonates  present.  It  has 
the  power  of  absorbing  coloring-matter  and  certain  salts,  e.g.  lead 
salts,  from  liquids.  The  charcoal  obtained  from  the  dry  distilla- 
tion of  sugar  is  noted  for  its  peculiar  lustre.  These  different 


177.] 


ALLOTROPIC  FORMS   OF  CARBON. 


245 


sorts  of  charcoal  do  not  consist  of  pure  carbon  but  contain  other 
substances  in  small  proportions.  It  is  a  general  rule  that  carbon 
conducts  heat  and  electricity  better  the  longer  it  has  been  exposed 
to  a  high  temperature. 

177.  The  various  kinds  of  carbon  all  find  their  respective  uses.  Soot, 
or  lampblack,  serves  for  the  preparation  of  India  ink  and  black  paint. 
Gas  carbon  (coke),  being  a  good  conductor  of  electricity,  is  used  in  the 
electrical  industry.  Wood  charcoal  is  used  in  the  manufacture  of  gun- 
powder ;  animal  charcoal,  or  bone-black,  as  a  water-filter  to  remove  color- 
ing-matter, ill-smelling  gases  or  injurious  salts  (lead  salts)  from  drink- 
ing-water ;  it  is  also  employed  in  enormous  quantities  in  sugar  refineries 
to  decolorize  sugar  liquids. 

By  far  the  most  important  use  of  carbon  is  as  a  f  u  e  1 .  The  heat 
generated  by  the  burning  of  coal  warms  our  houses,  drives  our  steam- 
engines,  etc. 

The  principal  kinds  used  as  fuels  are  charcoal,  coke,  anthracite  coal, 
bituminous  coal,  brown  coal  (lignite)  and  peat. 

Charcoal  (wood  charcoal)  is  made  on  a  large  scale  by  the  colliers. 
Long  sticks  of  wood  are  piled  in  a  large  heap,  covered  with  sod  and 
ignited  at  the  bottom.  The  wood  smolders  away  slowly  and  becomes 
completely  charred.  This  "  charcoal-pit "  process  is  not  at  all  economical, 
inasmuch  as  all  the  volatile  products  are  lost;  it  is  carried  on  exten- 
sively (Fig.  42),  but  it  is  being  more  and  more  replaced  by  the  dry  disr 
filiation  of  wood  from  iron  retorts,  in  which  process  the  gaseous  and 
tarry  products  are  recovered. 


FIG.  42. — CHARCOAL  PIT. 

Coke  is  the  residue  in  the  retorts  of  the  gas  factories  after  the  coal 
has  been  deprived  of  its  volatile  products  by  heating.  It  is  also  manu- 
factured on  a  large  scale  for  metallurgical  and  other  purposes.  Coke  is 
thought  by  many  to  have  a  great  future  as  a  fuel,  since  it  is  a  hard- 
burning  smokeless  fuel,  manufactured  from  the  cheap  soft  coal. 


246 


INORGANIC  CHEMISTRY. 


[§§  177- 


Peat  and  the  various  coals  owe  their  origin  to  the  same  geological 
process,  the  slow  decay  of  plant-remains  in  the  absence  of  air.  Peat  is 
the  youngest  formation  and  anthracite  coal  the  oldest.  During  this 
transition  carbon  dioxide  and  methane,  CH4,  are  given  off  and  the  residue 
becomes  richer  in  carbon  and  poorer  in  hydrogen  and  oxygen  than  the 
corresponding  chief  constituent  of  plant  tissues,  cellulose.  The  follow- 
ing table  shows  this: 


Carbon. 

Hydrogen. 

Oxygen. 

Cellulose  

50.0% 

6.0% 

44  0% 

Peat  

60.0 

5.9 

34  1 

Brown  coal  

67.0 

5.8 

27  2 

Cannel  coal 

85  8 

5  8 

8  3 

Anthracite  coal 

94  0 

3  4 

2  6 

The  plants  of  which  these  formations  originally  consisted  are  different. 
Peat  appears  from  its  structure  to  have  come  chiefly  from  swampy 
growths,  mosses  and  the  like;  mineral  coal  from  extinct  plants,  gigantic 
horsetails  (equiseta),  lepidodendra  and  sigillariae. 

Molecular  and  Atomic  Weight  of  Carbon. — Chemical  Properties. 

178.  The  carbon  molecule  probably  contains  a  large  number  of 
atoms.  It  has  not  yet  been  possible  to  determine  how  large  this 
number  is.  It  is  supposed  that  graphite  has  a  larger  number  of 
atoms  to  the  molecule  than  amorphous  carbon,  and  diamond  more 
than  graphite,  since  graphite  and  diamond  are  less  easily  attacked 
by  chemical  reagents  and  because  they  are  denser.  • 

A  determination  of  the  vapor  density  of  carbon  is  of  course 
out  of  the  question.  The  measurement  of  the  melting-point 
depression  that  carbon  produces  in  iron  is  also  impracticable; 
however,  it  is  known  that  even  a  small  percentage  of  carbon  causes 
a  considerable  lowering  of  the  melting-point  of  iron  (see  §  304). 

It  can  be  shown  in  the  following  way,  however,  that  the  number 
of  atoms  in  the  carbon  molecule  must  be  very  great.  By  the 
oxidation  of  amorphous  carbon  with  potassium  permanganate 
mellitic  acid  is  formed,  which  contains  12  carbon  atoms  to  the 
molecule.  This  makes  it  quite  probable  that  the  carbon  mole- 
cule contains  at  least  12  atoms,  for  in  the  oxidation  of  organic 
substances  the  products  almost  always  contain  either  a  smaller 
or  the  same  number  of  carbon  atoms  to  the  molecule.  For  the 


178.]    MOLECULAR  AND  ATOMIC  WEIGHT  OF  CARBON.        247 

following  reason  it  is,  however,  to  be  supposed  that  the  number 
of  atoms  in  the  carbon  molecule  is  much  greater  than  12.  When 
marsh-gas,  CH4,  is  passed  through  a  red-hot  tube,  ethylene,  C2-H4, 
is  formed  among  other  things.  If  this  is  then  treated  in  the  same 
way,  acetylene,  C2H2,  is  obtained,  and  from  this  again  benzene, 
C6H6.  On  conducting  benzene  vapor  through  a  glowing  tube, 
naphthalene,  CioH8,  pyrene,  Ci6H10,  etc.,  are  formed.  If  either 
of  the  latter  is  heated  still  higher  (in  the  absence  of  air)  carbon 
is  deposited.  We  thus  see  that  as  the  temperature  rises  the  num- 
ber of  carbon  atoms  in  the  molecule  steadily  increases.  The 
final  product  of  these  operations,  carbon,  will  therefore  probably 
contain  a  considerably  larger  number  of  atoms  in  its  molecule 
than  naphthalene  or  pyrene. 

Carbon  can  unite  directly  with  many  elements.  At  ordinary 
temperatures  it  combines  with  fluorine  only.  MOISSAN  intro- 
duced lampblack  into  fluorine  gas,  and  the  carbon  commenced 
to  glow;  when  fluorine  was  present  in  excess  carbon  tetraftuoride, 
CF4,  was  formed. 

Hydrogen  combines  with  carbon  directly  to  form  acety- 
lene and  a  small  quantity  of  marsh-gas,  when  an  electric  arc  is 
passed  between  two  carbons  in  an  atmosphere  of  hydrogen.  Of 
all  the  many  compounds  consisting  of  only  carbon  and  hydrogen 
these  are  the  only  ones  which  can  be  obtained  by  direct  synthesis. 
Under  analogous  conditions  carbon  unites  with  chlorine  to 
form  perchloroethane,  C2C16,  and  hexachlorobenzene,  CeCle- 

Oxygen  unites  with  carbon  at  an  elevated  temperature  to 
form  carbon  monoxide,  CO,  or  carbon  dioxide,  C02,  according  as 
carbon  or  oxygen  is  in  excess.  If  sulphur  vapor  is  passed  over 
red-hot  coals,  carbon  disulphide,  CS2,  is  produced. 

The  elements  of  the  nitrogen  group,  N,  P;  As,  Sb  and 
Bi,  do  not  combine  with  carbon  directly.  Silicon  and  car- 
bon unite  at  the  temperature  of  the  electric  furnace  to  form  CSi, 
carborundum,  which  is  so  hard  that  it  can  be  used  as  a  powder  for 
polishing  glass  and  precious  stones. 

MOISSAN  also  found  that  many  metals  are  able  to  combine 
with  carbon  at  a  very  high  temperature,  forming  carbides.  This 
was  previously  known  to  be  true  of  iron  and  certain  other 
metals. 


248  INORGANIC   CHEMISTRY.  [§§  179- 

The  difference  in  the  behavior  of  these  carbides  towards  water  is 
interesting.  Iron  carbide  is  unaffected  by  it;  calcium  carbide  gives 
acetylene,  C2H2;  aluminium  carbide  yields  methane;  other  carbides 
give  mixtures  of  the  two  hydrocarbons;  uranium  carbide  produces 
methane  and  also  liquid  and  solid  hydrocarbons. 

179.  The  atomic  weight  of  carbon  has  been  determined  with 
great    accuracy   by    DUMAS    and   STAS.     The    averages    for   the 
different   series    of   experiments,    each    of   which    showed   little 
variation,  were  as  follows: 

Ratio  by  weight  of  carbon  to  oxygen  in  carbon  di- 
oxide from  the  combustion  of: 

Natural  graphite 2.9994:8.0000 

Artificial       "       2.9995:8.0000 

Diamond 3.0002:8.0000 

The  ratio  of  carbon  to  oxygen  in  carbon  dioxide  is  thus  very 
close  to  3:8.  As  the  specific  gravity  of  carbon  dioxide  points  to 
a  molecular  weight  of  44  for  this  gas,  it  must  contain,  according 
to  this  ratio,  12.00  parts  by  weight  of  carbon  and  32  parts  of 
oxygen.  The  formula  is  therefore  G/^.  Inasmuch  as  no  carbon 
compound  is  known  whose  molecular  weight  includes  less  than  12 
parts  of  carbon,  we  have  CO2  as  the  formula;  hence  the  atomic 
weight  of  carbon  must  be  12.00  for  0=16. 

Compounds  with  Hydrogen. 

180.  Carbon  and  hydrogen  form  a  very  large  number  of  com- 
pounds (hydrocarbons),  which  are  more  fully  discussed  in  organic 
chemistry.    Two  of  them  will  be  treated  here  briefly. 

Methane,  also  called  marsh-gas  and  fire-damp,  is  the  only  hydro- 
carbon with  just  one  atom  of  carbon.  It  occurs  in  nature  in 
volcanic  gases;  moreover,  it  gushes  out  of  the  ground  in  the  neigh- 
borhood of  'the  oil-wells  at  Baku  and  various  places  in  America. 
It  is  an  important  constituent  of  "  natural  gas."  It  owes  the 
name  "marsh-gas"  to  the  fact  that  it  arises  from  swamps,  especially 
when  the  decaying  vegetation  at  the  bottom  is  stirred  up.  It 
is  called  "  fire-damp  "  because  it  occurs  in  coal  beds  (§  177),  from 
which  it  escapes  when  they  are  broken  up.  It  forms  a  violently 
explosive  mixture  with  air,  which  is  frequently  the  cause  of  mine 
explosions.  For  its  modes  of  formation  and  its  physical  and  chem- 
ical properties  reference  should  be  had  to  OBG.  CHEM.,  §  29. 


181.]  COMPOUNDS  WITH  OXYGEN.  249 

181.  Acetylene,  C2H2,  is  a  colorless  gas  of  a  disagreeable  odor.     It  is 
soluble  in  an  equal  volume  of  water  at  18°  and  becomes  liquid  at  18° 
under  83  atmospheres.     Its  hydrogen  atoms  are  replaceable  by  metals. 
It  is  manufactured  by  decomposing  calcium  carbide  with  water  : 
CaC2+2H20=Ca(OH)2  +  C2H2. 

Calcium  carbide  is  prepared  by  heating  coke  with  unslaked  lime 
(CaO)  in  the  electric  furnace.  The  calcium  formed  by  the  action  of 
carbon  on  lime  unites  with  carbon  at  the  high  temperature  of  the  fur- 
nace to  form  CaC2.  Acetylene  burns  with  a  vivid  flame  on  coming 
out  of  a  small  orifice  under  pressure.  Since  it  can  be  prepared  from 
calcium  carbide  pretty  cheaply,  it  is  used  rather  extensively  in  small 
systems  for  illuminating  purposes.  When  mixed  with  air  and  ignited 
it  explodes  vehemently;  the  compounds  with  metals  are  also  explosive. 
It  is  endothermic  and  can  be  exploded  by  fulminating  mercury. 

The  combustion  of  acetylene  is  another  illustration  of  the  rule  of 
§  137,  that  reactions  are  in  most  cases  of  a  simpler  nature  than  the 
chemical  equations  indicate.  The  equation  here  is: 

2C2H2+502=4C02+2H20. 

According  to  this  equation  the  combustion  should  be  septimolecular. 
BONE  and  CAIN  proved,  however,  that  the  reaction  has  more  than  one 
stage,  the  first  stage  being  represented  by  the  bimolecular  equation  : 


CO  and  H2  then  burn  further  to  CO2  and  H2O. 

From  a  kinetic  standpoint,  it  is  quite  conceivable  that  polymolecular 
reactions  should  be  rare,  for  the  probability  of  a  large  number  of  mole- 
cules coming  together  in  just  such  a  way  that  a  reaction  can  take  place 
is  indeed  very  slight.  The  reaction  is  more  likely  to  proceed  in  a  way 
which  involves  the  interaction  of  only  very  few  molecules. 

Compounds  with  Oxygen. 

Three    oxygen    compounds    of    carbon    are    known:     carbon 
monoxide,  CO,  carbon  dioxide,  CO2,  and  carbon  suboxide}  CsC2. 
For  the  latter  compound,  see  ORG.  CHEM.  §  166. 

CARBON   MONOXIDE,  CO. 

182.  This  gaseous  compound  is  always  formed  when  carbon 
burns  in  a  limited  supply  of  air  or  oxygen.  A  number  of  carbon 
compounds  also  yield  carbon  monoxide  when  burned  under  this 
same  condition.  It  can  also  be  obtained  by  the  action  of  carbon 
on  oxygen  compounds,  e.g.,  by  heating  zinc  oxide,  ZnO  with 


250  INORGANIC  CHEMISTRY.  [§§  i82_ 

carbon.     On  passing  steam  over  red-hot  coals  a  mixture  of  hydrogen 
and  carbon  monoxide  is  produced: 


This  mixture  goes  by  the  name  of  water-gas.  It  is  used  on  a  large 
scale  for  heating  and  lighting,  especially  in  America.  For  the  latter 
purpose  it  is  charged  with  the  vapor  of  hydrocarbons  rich  in  carbon, 
since  its  own  flame  is  not  luminous.  The  use  of  the  incandescent  gas- 
light (§291)  makes  this  "  carburetting  "  unnecessary.  Water-gas  con- 
taining 50%  of  carbon  monoxide  is  very  poisonous  (ORG.CHEM.  §  241). 

Carbon  monoxide  is  also  formed  by  the  reduction  of  carbon 
dioxide  with  red-hot  carbon: 

C  +  C02=2CO. 

This  reaction  is  limited  by  the  reverse  one  and  we  have  here  a 
case  of  balanced  action  expressed  by 


In  view  of  the  caloric  effect  of  the  reaction, 
2CO=C  +  CO2  +  3900CaL, 

an  elevation  of  temperature  must,  according  to  LE  CHATELIER'S 
rule  (§  51),  increase  the  amount  of  carbon  monoxide;  a  depression 
of  temperature,  the  opposite.  Experience  has  shown  this  to  be 
actually  the  case.  As  the  temperature  rises  the  quantity  of  carbon 
monoxide  increases  rapidly  and  at  1000°  there  is  still  a  very  small 
amount  of  dioxide.  At  445°,  on  the  other  hand,  practically  all 
the  carbon  monoxide  is  changed  into  carbon  dioxide  and  carbon. 
This  result  is  surprising,  because  the  same  change  should  also 
occur  at  lower  temperatures;  nevertheless,  carbon  monoxide  seems 
perfectly  stable  at  ordinary  temperatures,  even  as  high  as  200°. 
The  cause  of  this  phenomenon  must,  as  in  analogous  cases,  be 
sought  in  the  very  great  retardation  of  the  velocity  of  the  reaction 
2CO—  >CO2  +  C  when  the  temperature  sinks.  On  using  certain 
catalyzers,  e.g.  finely  divided  nickel,  the  velocity  of  the  reaction 
2CO—  >CO2  +  C  becomes  measurable  as  low  as  256°. 

These  measurements  have  shown  that  the  decomposition  of 
carbon  monoxide  into  carbon  dioxide  and  carbon  is  not  a  bimolecu- 
lar  reaction,  as  would  be  expected  from  the  above  equation,  but 


183j  CARBON  MONOXIDE.  251 

a  unimolecular  one.  To  explain  this  it  may  be  suggested  that 
the  decomposition  takes  place  in  two  stages:  I.  CO=C-fO; 
II.  CO  +  O=CO2-  If  we  assume  that  the  second  stage  has  an 
infinite  velocity,  it  is  only  the  first  that  is  really  measured,  i.e.  a 
unimolecular  reaction. 

The  reduction  of  salts  of  carbonic  acid  also  furnishes  a  method 
of  preparing  carbon  monoxide.  If  chalk  (CaCOs)  or  magnesite 
(MgCOs)  is  heated  with  zinc  dust,  pure  carbon  monoxide  is  formed: 

CaC03  +  Zn  =  CaO  +  ZnO  +  CO. 

Physical  Properties. — Carbon  monoxide  is  a  colorless,  odorless 
gas  of  a  specific  gravity  of  0.967  (air  =  l).  It  is  hard  to  condense, 
its  critical  temperature  being  —139.5°  and  its  critical  pressure 
35.5  atmospheres.  It  boils  at  -190°  and  solidifies  at  -211°.  It 
is  only  slightly  soluble  in  water. 

183.  Chemical  Properties. — Carbon  monoxide  burns  with  a 
characteristic  blue  flame  to  carbon  dioxide.  It  can  unite  with 
chlorine  directly  to  form  phosgene,  COC12,  and  also  with  sulphur 
(at  an  elevated  temperature)  to  form  carbon  oxy  sulphide,  COS? 
both  compounds  are  gaseous.  Again,  it  unites  directly  with 
nickel  and  iron,  giving  the  compounds  Ni(CO)4  and  Fe(CO)5 
(§§  214  and  311). 

On  account  of  its  tendency  to  combine  with  oxygen,  it  displays 
strong  reducing  power,  especially  at  high  temperatures.  Thus 
metallic  oxides,  like  Fe20s,  CuO,  etc.,  are  easily  converted  into  the 
metals  when  hot.  Some  compounds  are  reduced  by  carbon  mon- 
oxide even  at  ordinary  temperatures.  Palladium  is  precipitated 
from  an  aqueous  solution  of  palladious  chloride  and  an  ammoniacal 
silver  solution  (prepared  by  dissolving  silver  oxide  in  ammonium 
hydroxide  to  the  point  of  saturation)  is  turned  black  by  the  gas 
on  account  of  formation  of  the  metal.  Both  of  these  reactions  serve 
for  the  detection  of  carbon  monoxide. 

An  ammoniacal  cuprous  chloride  solution  absorbs  the  gas 
because  of  the  formation  of  a  compound,  Cu2Cl2-CO  +  2H2O, 
which  can  be  isolated  in  the  crystalline  state  but  decomposes  again 
very  readily. 

The  composition  of  carbon  monoxide  can  be  determined  by 
exploding  a  mixture  of  the  gas  with  oxygen.  It  is  then  found  that 
2  vols.  CO  unite  with  1  vol.  O2  to  form  2  vols.  CO2.  This  together 
with  the  vapor  density  establishes  the  formula  as  CO. 


252  INORGANIC   CHEMISTRY.  [§§  183- 

CARBON   DIOXIDE,   CARBONIC  ACID   ANHYDRIDE,   C02. 

184.  This  compound  occurs  not  only  by  itself  but  also  in  com- 
bination. It  is  a  regular  constituent  of  the  air  (§  106);  many 
mineral  waters  contain  the  free  gas;  in  some  places  of  the  earth 
(in  the  Dog's  Grotto  at  Naples  and  the  famous  Poison  Valley  in 
Java)  it  comes  up  out  of  the  ground  and  it  is  also  found  in  volcanic 
exhalations.  The  most  minerals  and  rocks  contain  numerous 
extremely  small  cavities,  partly  filled  with  liquid  carbon  dioxide. 
Combined,  it  occurs  in  large  quantity  in  the  carbonates  of  lime 
and  magnesia  (§  176). 

Carbon  dioxide  results  from  the  combustion  of  carbon  in  an 
excess  of  oxygen  and  also  from  the  direct  decomposition  of  many 
salts  of  carbonic  acid  (carbonates)  by  •  heat  : 


2NaHCp3=Na2C03+H2O  +  CO2;  CaC03=CaO  +  C02. 

Sodium  bi- 
carbonate. 

Moreover,  it  is  formed  when  a  carbonate  is  decomposed  by  an 
acid: 

Na2C03  +  2HC1  =  2NaCl  +  H20  +  C02. 

By  the  action  of  oxygen  at  high  temperatures  all  carbon  com- 
pounds are  burned  with  the  formation  of  carbon  dioxide.  It  is 
also  produced  by  the  action  of  carbon  on  oxygen  compounds,  e.g. 
by  heating  powdered  charcoal  with  an  excess  of  copper  oxide; 
finally  also  by  the  interaction  of  carbon  compounds  and  oxygen 
compounds.  This  latter  action  is  the  basis  of  the  general  method 
for  determining  the  proportion  of  carbon  in  organic  substances; 
they  are  heated  together  with  copper  oxide  and  the  carbon  dioxide 
formed  is  absorbed  in  a  weighed  amount  of  caustic  potash. 

Physical  Properties.  —  Carbon  dioxide  at  ordinary  temperatures 
and  pressures  is  a  gas  with  a  somewhat  pungent  odor  and  taste. 
gp.  g.  =  1.529  (air=l).  It  is  thus  about  half  again  as  heavy  as 
air,  so  that  in  those  places  where  it  comes  out  of  the  earth,  as  in 
the  Dog's  Grotto  at  Naples,  it  stays  in  a  layer  close  to  the  ground 
and  a  dog,  for  instance,  is  suffocated  while  a  man  can  breathe 
with  comfort.  Carbon  dioxide  is  easily  condensed,  becoming  liquid 
at  0°  under  35  atmospheres  pressure.  Its  critical  temperature  is 
31.35°  and  its  critical  pressure  72.9  atm.  Liquid  carbon  dioxide 
("  liquid  carbonic  acid  ")  is  manufactured  in  great  quantities  and 


184.]  CARBON  DIOXIDE.  253 

brought  on  to  the  market  in  steel  bottles  (bombs).  It  is  a  very 
mobile  liquid,  which  is  not  miscible  with  water  in  all  proportions. 
If  the  liquid  is  allowed  to  escape  from  the  bomb  into  a  coarse 
linen  bag  (by  inverting  the  bomb  and  opening  the  valve),  part  of 
it  vaporizes,  absorbing  hereby  so  much  heat  that  the  remainder 
solidifies  in  white  flakes.  A  mixture  of  this  solid  carbon  dioxide 
with  ether,  alcohol  or  acetone  is  often  used  as  a  freezing-mixture; 
it  enables  us  to  obtain  a  temperature  of  —80°,  and  even  —140° 
in  vacua.  When  liquid  carbon  dioxide  is  cooled  down  in  a  sealed 
tube,  it  congeals  to  an  icy  mass,  which  melts  at  -65°. 

At  15°  carbonic  acid  gas  dissolves  in  its  own  volume  of  water 
(more  accurately  1.0020  vol.);  at  0°  in  1.7967  vol.  In  alcohol  it 
is  still  more  soluble. 

Chemical  Properties. — Carbon  dioxide  is  a  very  stable  com- 
pound; it  is  only  decomposed  by  intense  heat  (see  §  182)  or  by  the 
continued  action  of  induction  sparks,  breaking  up  into  oxygen  and 
carbon  monoxide.  This  decomposition  never  completes  itself,  for 
just  so  soon  as  a  certain  amount  of  these  gases  have  been  formed, 
they  reunite  with  explosion.  At  the  moment  before  the  explosion 
the  amount  of  carbon  dioxide  still  present  becomes  no  longer  suffi- 
cient to  dilute  the  mixture  of  oxygen  and  monoxide  enough  to 
hinder  an  explosion;  the  explosive  limit  is  reached. 

Carbon  dioxide  cannot  be  farther  oxidized;  it  is  therefore  not 
combustible.  In  general  it  cannot  support  combustion  either. 
There  are,  however,  certain  substances  that  take  up  oxygen  from 
it  when  hot;  if  carbon  dioxide  is  mixed  with  hydrogen  and  passed 
through  a  red-hot  tube,  carbon  monoxide  and  water  are  formed; 
when  led  over  glowing  carbon  or  when  heated  with  phosphorus  it 
is  reduced  to  carbon  monoxide.  If  a  burning  magnesium  ribbon 
is  lowered  into  carbon  dioxide,  the  oxide  of  the  metal  is  formed 
and  free  carbon  is  deposited ;  the  same  thing  happens  when  sodium 
or  potassium  is  heated  in  dry  carbon  dioxide. 

The  aqueous  solution  of  carbon  dioxide  reacts  slightly  acid;  it 
is  supposed  that  this  solution  contains  a  compound  H2CO3,  of 
which  many  salts  are  known.  This  acid,  carbonic  acid,  has  not 
yet  been  isolated  in  the  free  state,  however,  since  it  gives  off  gaseous 
carbon  dioxide  ("carbonic  acid  gas")  when  its  solution  is  boiled 
or  frozen.  If  its  salts  (carbonates)  are  treated  with  an  acid,  no 
H2C03  is  obtained  either,  for  it  breaks  up  forthwith  into  water 
and  carbon  dioxide.  Carbonic  acid  is  a  very  weak  acid;  it  is 


254  INORGANIC  CHEMISTRY.  [§§  184- 

liberated  from  its  salts  by  almost  every  other  acid.  By  adding 
hydrochloric  acid  to  a  carbonate  H*  ions  are  introduced  into  the 
liquid  and  they  unite  with  the  COg"  ions  to  form  integral  H2COs 
molecules.  These,  however,  break  up  largely  into  water  and 
carbon  dioxide,  the  latter  of  which  can  only  remain  in  solution 
up  to  a  certain  amount  at  a  constant  pressure,  so  that  all  in  excess 
of  this  passes  out.  As  a  result  the  concentration  of  the  H2COs 
molecules  cannot  exceed  a  definite  and,  in  this  case  low  limit. 
Since,  however,  the  ionization  of  these  molecules  is  very  weak, 
in  reality  all  of  the  carbonate  is  decomposed  by  the  strong  acid 
(§73).  ' 

The  neutral  carbonates  of  the  alkalies  are  soluble  in  water, 
giving  an  alkaline  reaction,  as  a  result  of  partial  hydrolysis  (§  66). 
The  acid,  H2CO3,  is  a  weak  acid,  although  its  salts,  e.g.  K2CO3, 
are  strong  electrolytes.  A  solution  of  such  a  salt,  therefore,  con- 
tains a  large  number  of  COs"  ions,  part  of  which  must  unite  with 
the  H*  ions  of  the  water  in  order  to  establish  the  equilibrium  be- 
tween carbonic  acid  and  its  ions.  The  result  of  this  is  that  other 
molecules  of  water  must  be  split  up  into  ions  in  order  to  com- 
pensate the  loss  of  H*  ions.  This  leaves  in  the  liquid  a  certain 
number  of  OH'  ions,  which  are  not  balanced  by  an  equal  number 
of  H"  ions.  The  liquid  therefore  acquires  an  alkaline  reaction. 

The  carbonates  of  the  other  metals  are  insoluble  in  water; 
however,  the  acid  carbonates  are  mostly  soluble.  Calcium  carbo- 
nate, e.g.,  dissolves  in  water  containing  carbonic  acid.  The  solu- 
tions of  such  acid  carbonates  give  off  carbon  dioxide  on  merely 
boiling,  however,  and  the  neutral  carbonates  are  precipitated. 
In  the  solid  state  also  the  acid  carbonates  give  off  carbonic  acid 
gas  very  readily  on  warming. 

Composition  of  Carbon  Dioxide. — In  connection  with  what  was 
stated  in  §  179  it  is  an  important  fact  that  no  change  of  volume 
occurs  when  carbon  burns  in  an  excess  of  oxygen: 

C+O2=CO2. 

1  vol.      1  vol. 

When  a  very  concentrated  solution  of  potassium  carbonate  is  elec- 
trolyzed  with  high  current  density  at  30°-40°,  potassium  percarbonate, 
K2C2O6,  is  formed  at  the  anode.  In  aqueous  solution  it  sets  free  iodine 
from  KI  solution  at  once,  which  serves  to  distinguish  it  from  H202,  since 
a  dilute  solution  of  the  latter  liberates  iodine  only  very  slowly. 


185.]  OTHER  CARBON  COMPOUNDS.  255 

Other  Carbon  Compounds. 

185.  Cyanogen,  (CN)2,  can  be  prepared  by  heating  mercuric 
cyanide,  Hg(CN)2,  or  by  treating  a  solution  of  potassium  cyanide 
with  copper  sulphate  solution.  It  is  possible  that  first  cupric 
cyanide  is  formed  and  that  this  at  once  breaks  up  into  cuprous 
cyanide  and  cyanogen: 


Cyanogen  has  a  penetrating  odor.  When  liquefied  it  boils  at 
—  20.7°.  It  is  unaffected  by  high  temperatures.  It  dissolves  in 
water,  but  the  solution  deposits  amorphous  brown  flakes  after  a 
while.  It  burns  with  a  purple-,tinged  flame  according  to  the 
equation 


The  reaction,  however,  is  not  trimolecular,  the  first  stage  being 


i.e.  a  bimolecular  process. 

This  was  proved  by  DIXON  by  determining  the  velocity  of  propagation 
of  the  explosion  of  mixtures  of  cyanogen  and  oxygen.  When  explosive 
gas  mixtures  are  introduced  into  a  long  tube  and  their  explosion  started 
at  one  end  (by  an  electric  spark,  for  example)  a  flame  results,  which  is 
propagated  through  the  tube  with  a  definite  and  measurable  velocity. 
BERTHELOT  called  this  self-propagating  flame  the  explosion  wave. 

DIXON  ignited  a  mixture  of  1  vol.  cyanogen  and  1  vol.  oxygen,  obtaining 
after  the  explosion  carbon  monoxide  and  nitrogen;  the  velocity  of  the 
explosion  wave  was  found  to  be  2728  m.  per  sec.  Thereupon  he  mixed  1  vol. 
cyanogen  with  2  vols.  oxygen  in  one  instance  and  with  1  vol.  oxygen  and 
1  vol.  of  an  indifferent  gas  in  another  instance;  in  both  cases  the  velocity 
of  the  explosion  wave  was  nearly  the  same,  viz.  2321  m.  and  2398  m.  per 
sec.  It  is  plain,  therefore,  that  the  second  volume  of  oxygen  influenced 
the  explosion  wave  in  the  same  way  as  the  indifferent  gas,  viz.  as  a  diluent. 
The  conclusion  may  be  drawn  that  in  the  explosion  wave  itself  only  carbon 
monoxide  and  nitrogen  are  formed,  even  in  the  presence  of  an  excess  of 
oxygen.  However,  since  the  tube  contains  only  carbon  dioxide  and  nitrogen 
after  the  combustion,  it  must  be  assumed  that  the  combustion  of  carbon 
monoxide  to  carbon  dioxide  is  a  secondary  process. 


256 


INORGANIC   CHEMISTRY. 


185- 


Hydrogen  cyanide,  HCN  (prussic  acid),  is  important  in  inor- 
ganic chemistry  because  of  the  numerous  complex  salts  which  it 
forms.  Those  of  the  alkalies  are  soluble  in  water  and  crystallize 
beautifully;  see  §  308.  The  salts  of  the  alkaline  earths  and 
mercuric  cyanide  are  also  soluble  in  water,  the  other  salts  in- 
soluble. 

The  Flame. 

186.  A  flame  is  produced  by  the  burning  of  a  gas;  solids,  like 
iron,  carbon,  etc.,  burn  without  a  flame.  If  a  flame  is  observed 
during  the  burning  of  mineral  coal,  a  candle  or  the  like,  it  is  due 


FIG.  43. — HEVERSE  FLAME.       FIG.  44. — POTASSIUM  CHLORATE  FLAME. 

to  the  fact  that  at  that  high  temperature  gaseous  decomposition- 
products  are  formed,  which  burn.  If  a  gas  burns  in  the  air,  it  is 
called  a  combustible  gas  and  the  oxygen  of  the  air  is  called  the  sup- 
porter of  the  combustion.  These  expressions  in  common  use  are 
only  relative  terms;  it  is  possible  to  light  the  oxygen  and  have  it 
burn  with  a  flame  in  a  gas  which  is  ordinarily  called  combustible. 
This  phenomenon  is  illustrated  in  a  way  by  the  reverse  flame. 

This  can  be  easily  obtained  with  the  aid  of  the  apparatus  of  Fig.  43. 
A  lamp-chimney  is  fitted  with  a  two-hole  cork  at  its  lower  end.     Through 


187-]  THE  FLAME.  257 

the  narrower  hole  of  the  cork  a  small  tube  a  is  inserted  for  conducting 
in  the  gas;  through  the  wider  hole  a  tube  b  for  the  admission  of  air. 
The  chimney  is  first  removed  and  the  gas  coming  out  of  tube  a  lighted 
and  so  regulated  as  to  produce  a  small  flame.  Then  the  chimney  is 
replaced;  the  flame  continues  to  burn  quietly,  inasmuch  as  plenty  of 
air  io  supplied  by  the  wider  tube.  Thereupon  the  gas  supply  is  gradu- 
ally increased  and  at  a  certain  moment  the  small  flame  at  the  end  of  a 
is  extinguished  and  a  large  pale  flame  appears  at  the  end  of  b;  it  is  air 
burning  in  the  gas  which  fills  the  chimney.  This  is  the  reverse  flame 
of  air  in  illuminating-gas.  At  the  same  time  the  excess  of  gas  escaping 
at  the  top  ignites  in  the  outside  air,  so  that  the  apparatus  presents  both 
a  direct  and  a  reverse  flame  at  the  same  time.  That  it  is  really  air 
that  burns  at  the  mouth  of  b  is  proved  by  introducing  a  tiny  gas-flame 
by  means  of  the  tube  c  into  the  flame  of  the  wide  tube  6;  the  small 
flame  continues  to  burn. 

Substances  that  give  up  oxygen  are  capable  of  burning  when  sur- 
rounded by  a  combustible  gas.  The  experiment  can  be  carried  out 
with  potassium  chlorate  as  follows:  Illuminating-gas  is  conducted  into 
a  glass  cylinder  (Fig.  44)  and  lighted  at  the  top,  where  the  cylinder 
is  covered  by  a  thin  piece  of  metal  with  a  hole  in  it.  A  little  potassium 
chlorate  is  then  lowered  into  the  flame  by  means  of  a  deflagrating  spoon 
and  heated  till  oxygen  comes  off  freely.  If  the  bowl  is  then  dipped 
down  in  the  cylinder,  the  oxygen  burns  with  a  very  luminous  flame, 
which  is  colored  violet-blue  by  the  vaporization  of  some  potassium  salt. 

We  saw  above  (§27)  that  a  hydrogen  flame  continues  to  burn  in  chlo- 
rine with  the  formation  of  hydrochloric  acid;  on  the  other  hand  chlorine 
can  also  be  made  to  bum  in  hydrogen.  For  this  purpose  a  cylinder 
closed  at  the  top  is  filled  with  hydrogen  and  lit  at  the  lower  edge.  A 
tube  through  which  chlorine  is  supplied  is  then  brought  in  contact  with 
this  flame  and  inserted  in  the  cylinder.  The  chlorine  burns  on. 

187.  A  flame  may  be  luminous  or  non-luminous.  It  gives 
light  when  solid  particles  are  suspended  in  it.  An  ordinary  gas- 
flame  is  luminous  because  particles  of  carbon,  set  free  by  the  com- 
bustion, are  made  to  glow.  On  introducing  a  cold  object  into  the 
flame  they  are  deposited  as  soot.  The  light  of  the  WELSBACH 
incandescent  gas-light  is  produced  by  the  glowing  incombustible 
mantle  (§  291). 

Such  flames  give  a  continous  spectrum  (§  263).  Many 
gases,  which  yield  only  gaseous  products  on  burning,  give  either  a 
very  faint  light  or  none  at  all,  e.g.  hydrogen,  carbon  monoxide, 
etc.  However,  when  hydrogen  burns  in  oxygen  of  20  atmospheres 


258 


INORGANIC  CHEMISTRY. 


[§§  IBS- 


pressure,  its  flame  is  strongly  luminous.  Other  incandescent  gases, 
such  as  the  vapors  of  certain  metals;  can  render  a  flame  luminous 
even  at  ordinary  pressure,  imparting  to  it  a  definite  color.  Colored 
flames  of  this  sort  give  aline  spectrum  (§  263). 

A  gas-flame,  whose  light  is  due  to  incandescent  particles  of  car- 
bon, is  made  non-luminous  by  mixing  the  gas  with  air  before  the 
combustion.  This  is  the  principle  of  the  BUNSEN  burner  (Fig.  45), 
which  is  used  in  all  laboratories  and  quite  extensively  also,  with  some 
variation  or  other,  in  heating  and  cooking  apparatuses  (gas  stoves'). 

The  Bunsen  burner  consists  of  a  base  in 
which  is  a  tube  for  supplying  the  gas,  which 
escapes  from  a  narrow  orifice  at  a.  Here 
it  mixes  with  air  that  enters  through  the 
lateral  holes  in  c,  the  proportion  of  air  being 
regulated  by  the  collar  b.  This  mixture 
burns  with  a  colorless  flame  when  ignited 
at  the  top  of  c. 

The  opinion  was  originally  held  that  the 
loss  of  luminosity  of  the  flame  is  due  to 
the  oxygen  of  the  air,  the  latter  causing  the 
complete  combustion  of  the  carbon  particles. 
As  has  since  been  shown,  however,  the 
dilution  of  the  burning  gas  with  nitrogen  also 
has  a  part  in  it :  if  illuminating-gas  is  mixed 
with  two  or  three  times  as  much  nitrogen,  ^^~~ — 
the  former  burns  with  a  colorless  flame.  JTIG>  45^ BUNSEN  BURNER. 

With  the  aid  of  a  wire  gauze  a  burning  gas  mixture  can  be 
cooled  so  low  that  the  combustion  cannot  propagate  itself  through 
the  gauze;  in  other  words,  the  flame  does  not  get  through  the 
gauze  (Fig.  46).  If  gas  is  allowed  to  flow  out  of  a  BUNSEN  burner 
and  a  wire  gauze  is  held  across  the  current  a  short  distance  from 
the  orifice,  the  gas  can  be  lit  above  the  gauze  without  the  flame 
springing  back  to  the  burner. 

It  was  by  experiments  such  as  these  that  DAVY  was  led  to  discover 
his  miner's  safety -lamp.  As  Fig.  47  shows,  this  consists  of  an  oil-lamp, 
the  flame  of  which  is  surrounded  by  a  wire  cage.  A  combustible  gas 
mixture  may  catch  fire  inside  of  the  lantern,  but  the  fire  cannot  pass 
through  the  gauze  to  the  outside. 


188-] 


THE  FLAME. 


259 


188.  The  temperature  of  the  flame  is  much  lower  than  we  might 
suppose.  Since,  when  hydrogen  burns  in  oxygen,  57.2  kg.-calories 
are  produced  by  every  18  g.  of  the  mixture,  and  the  specific  heat  of 


FIG.  46. — EFFECT  OF  A  WIRE  GAUZE  ON  A  FLAME. 

steam  is  0.48,  this  amount  of  heat  ought  to  raise  the  18  g.  steam  to 

57  2 

a  temperature  of  ^         '       6"66QO°.     In  reality  the  temperature 

U.UloX  U  4o 

of  the  flame  does  not  exceed  2500°.  This  difference  between 
calculation  and  observation  is  due  to  the  fact  that  on  account 
of  dissociation  only  a  partial  combination  of  hy- 
drogen and  oxygen  takes  place  in  any  part  of 
the  flame.  The  temperature  of  6600°  could  indeed 
be  obtained  at  any  point,  if  the  gases  united  there 
completely  and  instantaneously;  but  this  is  im- 
possible, for  above  1300°  the  formation  of  the  com- 
pound is  checked  by  the  opposite  process,  the  dis- 
sociation of  steam.  Therefore  what  occurs  must 
be  this:  oxygen  and  hydrogen,  when  brought  to- 
gether at  the  aperture,  combine  and  effect  a  certain 
rise  of  temperature;  in  the  same  measure  as  the 
system  in  equilibrium  (hydrogen,  oxygen,  steam) 
FIG.  47.— MINER'S  cools  off,  fresh  portions  of  the  gases  unite.  Their 
SAFETY-LAMP,  combustion  cannot  therefore  take  place  at  any 
particular  point  but  must  be  gradual  throughout  the  whole  extent 
of  the  flame  and  at  any  one  point  the  temperature  cannot  ex- 
ceed a  certain  limit,  which  is  determined  by  the  degree  of  dissocia- 
tion of  the  combustion  product. 


260  INORGANIC    CHEMISTRY.  [§§  189- 

189.  Zones  oj  a  luminous  flame.     Let  us  take  a  candle-flame,  for 
example.     In  the  central  zone   (1  in  the  diagram  Fig.   48) 


FIG.  48.— ZONES  OF  A  LUMINOUS  FLAME. 

there  is  no  combustion.  The  stearin  of  the  candle  is  here  con- 
verted by  the  heat  of  the  flame  into  volatile  combustible  products. 
In  a  large  candle  this  can  be  proved  in  the  manner  shown  in  Fig. 
48.  The  narrow  tube  conducts  off  the  inflammable  gases  and  they 
can  be  lit  at  the  outer  end. 

The  hollo wness  of  a  flame  can  be  demonstrated  in  various  ways; 
in  a  Bunsen  burner,  for  instance,  by  placing  a  match-head  in  the 
center,  where  it  does  not  ignite,  or  by  holding  a  thin  platinum  wire 
across  a  flame;  the  wire  only  glows  at  the  edges  of  the  flame. 

The  dark  central  zone  of  the  flame  is  next  surrounded  by  the 
luminous  zone  (2) .  Here  the  volatilized  hydrocarbon  is  decom- 
posed with  the  separation  of  carbon,  because  the  air  supply  is 
insufficient  for  complete  combustion.  These  carbon  particles 
become  incandescent  and  so  make  the  flame  luminous.  Finally 
there  is  the  blue  outer  zone  (3),  in  which  the  glowing  particles 
of  carbon  are  burned  by  direct  contact  with  the  air.  It  radiates 
very  little  light. 

The  amount  of  solid  carbon  in  a  flame  which  is  raised  to  incandes- 
cence and  hence  gives  light  is  very  small,  as  the  following  calculation 
shows.  The  substances  in  burning  illuminating-gas  which  break  up 
with  the  liberation  of  carbon  are  chiefly  benzene  and  ethylene.  The 
former  makes  up  about  1,  the  latter  about  4,  per  cent  by  volume  of 
the  gas.  If  we  assume  that  the  benzene  is  completely  broken  up  and 


190.]  SILICON.  261 

the  ethylene  only  half,  then  the  total  amount  of  carbon  deposited  by 
1  liter  of  burning  illuminating-gas  is  about  54  mg.  The  volume  of 
the  luminous  part  of  a  gas  flame  with  a  consumption  of  150  liters  per 
hour  amounts  to  about  2  c.c.  (reduced  to  0°),  so  that  the  mass  of  solid 

2x54 
incandescent  carbon  present  in  it  is  only  =0.1  mg. 


SILICON. 

190.  This  element  in  combination  with  oxygen  is  one  of  the 
principal  constituents  of  the  earth's  crust  (§  8).  In  the  free  state, 
however,  it  does  not  occur  in  nature,  being  found  almost  exclusively 
as  silica,  SiO2,  or  in  the  silicates.  Sand  and  the  many  varieties  of 
quartz  are  different  forms  of  natural  silicon  dioxide;  the  number  of 
silicates  is  very  large.  . 

Free  silicon  is  obtained  by  heating  sodium  fluosilicate,  Na2SiFg, 
with  sodium: 

Na2SiF6  +  4Na  =  6NaF  +  Si, 

or  by  heating  sodium  in  an  atmosphere  of  silicon  tetrafluoride: 
4Na+SiF4=4NaF+Si. 

The  sodium  fluoride  can  be  removed  by  water. 

Another  method,  which  is  far  easier,  is  to  mix  400  g.  aluminium 
filings  with  500  g.  sulphur  and  360  g.  sand.  This  mixture  is  ignited, 
whereupon  it  burns  with  a  large  flame.  The  mass  fuses  and  becomes 
white-hot.  When  cooled  it  consists  principally  of  aluminium 
sulphide  and  free  silicon.  It  is  then  treated  with  dilute  hydro- 
chloric acid,  which  decomposes  and  dissolves  the  sulphide,  leaving 
the  silicon  behind.  (KUHNE  method.) 

Allotropic  Forms.-  —  The  silicon  obtained  by  the  two  first-named 
methods  is  a  brown  amorphous  powder;  it  can  be  fused 
under  a  layer  of  molten  sodium  chloride  and  obtained  crystal- 
line on  cooling.  The  latter  form  is  best  prepared  by  KUHNE'S 
method.  The  crystals  are  regular,  black,  and  of  a  high  lustre. 
If  silicon  is  heated  in  the  electric  furnace,  it  vaporizes  and  con- 
denses again  in  small  globules,  mixed  with  a  little  gray  powder 
and  some  silica. 

Chemical  Properties.  —  Silicon  takes  fire  only  when  heated  in 


262  INORGANIC  CHEMISTRY.  [§§  190- 

the  air  to  a  very  high  temperature,  burning  to  silica.  It  unites 
with  fluorine  at  ordinary  temperatures,  the  combustion  being 
marked  by  a  glow;  combination  with  chlorine  takes  place  on  gently 
warming.  At  an  elevated  temperature  silicon  combines  with  nitro- 
gen and  some  metals ;  these  silicides  have  been  prepared  mainly 
by  MOISSAN  in  his  electric  furnace. 

It  is  indifferent  towards  sulphuric,  nitric  and  hydrochloric 
acids.  Hydrofluoric  acid  dissolves  it,  however,  with  the  evolution  of 
hydrogen.  Hydrogen  chloride  gas  reacts  with  it  at  a  high  tem- 
perature, forming  silicon  tetrachloride  and  sili co-chloroform.  It 
dissolves  in  a  hot  solution  of  sodium  or  potassium  hydroxide,  pro- 
ducing hydrogen  and  a  silicate: 

Si + 2KOH + H20 = K2Si03 + 2H2. 

Hydrogen  Silicide,  SiH4. 

191.  This  gas  is  obtained  by  adding  freshly  prepared  magnesium 
silicide  to  hydrochloric  acid.  The  magnesium  silicide  is  prepared 
by  heating  sand  with  an  excess  of  magnesium  powder,  or  better  by 
fusing  40  parts  of  anhydrous  magnesium  chloride  with  a  mixture 
of  35  parts  of  sodium  fluosilicate,  10  of  sodium  chloride  and  20  of 
sodium.  The  hydrogen  silicide  so  obtained  is  mixed  with  hydro- 
gen. A  purer  product  results  from  heating  an  organic  derivative 
of  silicon,  tri-ethyl  silicof ormate : 

4SiH(OC2H5)3  =  3Si(OC2H5)4+SiH4. 

Hydrogen  silicide,  or  silicon  tetrahydride,  is  a  gas,  which 
becomes  liquid  at  —1°  under  a  pressure  of  100  atmospheres.  It  has 
a  disagreeable  odor.  It  takes  fire  in  the  air;  each  bubble  that 
escapes  from  the  generator  forms  a  cloudy  ring  of  hydrated  silica. 
If,  however  the  hydrogen  silicide  is  perfectly  pure,  it  does  not 
ignite  spontaneously  in  the  air  at  ordinary  temperatures  except 
under  reduced  pressure.  "  The  spontaneous  ignition  in  the  air 
is  caused  by  the  presence  of  small  quantities  of  other  compounds, 
probably  also  composed  of  silicon  and  hydrogen.  We  have 
therefore  in  this  case  phenomena  similar  to  those  in  the  case  of 
hydrogen  phosphide  (§  136).  Heat  decomposes  the  hydrogen 
silicide  readily  into  Si  and  2H2.  It  burns  in  a  chlorine  atmosphere 
and  is  decomposed  by  an  alkali  solution  according  to  the  equation : 

SiH4  +  2KOH + H20  =  4H2  +  K2SiO3. 


192]  HALOGEN  COMPOUNDS  OF  SILICON.  263 

Silico-ethane,  Si2H6,  is  formed  by  the  decomposition  of  magnesium 
silicide  by  hydrochloric  acid.  It  is  a  gas,  which  can  be  liquefied  below 
—  7° 

Halogen  Compounds  of  Silicon. 

192.  Silicon  tetrachloride,  SiCl4,  is  prepared  by  heating  silicon 
in  a  current  of  chlorine  at  300°-310°.  It  is  a  colorless  liquid  with 
the  specific  gravity  1.5241  at  0°  and  the  boiling-point  59.6°.  It 
is  instantly  decomposed  by  water,  forming  hydrochloric  acid  and 
hydrated  silica. 

Silico-chloroform,  SiCl3H,  is  obtained,  together  with  a  large  quantity 
of  silicon  tetrachloride,  on  heating  silicon  in  a  current  of  hydrochloric 
acid  gas  (§  190).  From  this  mixture  it  is  separated  by  fractional  distilla- 
tion. It  is  a  colorless,  strongly  smelling  compound  which  fumes  in 
the  air,  boils  at  34°,  and  is  decomposed  by  water. 

By  the  action  of  dark  electrical  discharges  on  a  mixture  of  dry 
hydrogen  and  silico-chloroform  vapors  chlorine-silicon  compounds  are 
formed  of  the  order  SinCl27i+2,  e.g.,  perchloro-silico-ethane,  Si2Cl6,  etc. 

Silicon  tetrafluoride,  SiF4,  can  be  obtained  by  warming  a 
mixture  of  sand  and  calcium  fluoride  with  concentrated  sulphuric 
acid: 

2CaF2  +  Si02 + 2H2S04  =  SiF4 + 2CaS04 + 2H20. 

It  is  a  colorless  gas  with  a  very  pungent  and  suffocating  odor; 
it  condenses  under  9  atm.  pressure  or  by  cooling  to  —160°.  When 
perfectly  dry,  it  does  not  attack  glass. 

Silicon  fluoride  is  also  formed  by  the  action  of  hydrogen  fluoride 
on  silicates ;  the  silica  is  first  set  free  from  them  and  then  attacked 
in  the  way  just  described.  Glass-etching  (§  53)  depends  on  this 
action. 

By  the  repeated  treatment  of  silicates  with  hydrous  hydrofluon 
acid  all  the  silicic  acid  is  driven  off  as  silicon  fluoride.  The  bases  which 
were  in  combination  with  the  silicic  acid  are  left  behind  in  the  form 
of  fluorides.  They  can  be  transformed  into  sulphates  by  warming 
with  sulphuric  acid  and  then  converted  into  a  form  suitable  for  analysis. 
We  have  here  a  very  useful  means  of  determining  the  metals  present  in 
the  silicates. 

Water  decomposes  silicon  fluoride  as  follows: 
3SiF4 + 3H20  =  H2Si03  +  2H2SiF6. 


INORGANIC  CHEMISTRY. 


[§§  192- 


The  compound  H2SiF6  is  called  hydrofluosilicic  acid;  it  is 
known  only  in  aqueous  solution.  If  the  latter  is  concentrated  by 
evaporation,  silicon  tetrafluoride  escapes  but  hydrogen  fluoride 
stays  in  solution.  When  the  concentration  corresponds  to  13.3% 
H^SiFe  the  vapor  contains  2HF  to  1SLF4;  but  dilute  solutions 
yield  a  vapor  which  contains  much  more  hydrogen  fluoride.  If, 
therefore,  a  concentrated  solution  of  hydrofluosilicic  acid  is  par- 
tially evaporated,  the  residual  liquid  is  able  to  dissolve  silica  because 
of  the  presence  of  free  hydrofluoric  acid.  On  the  other  hand, 
a  dilute  solution,  after  partial  evaporation,  leaves  a  residue,  from 
which  silicic  acid  is  deposited,  because  the  excess  of  silicon  tetra- 
fluoride which  it  contains  is  decomposed  by  water  according  to 
the  above  equation. 

The  decomposition  of  silicon  fluoride  by  water  is  usually  demon- 
strated in  the  following  way:  The  compound  is  generated  in  the  pre- 
scribed manner  in  a  flask  (Fig.  49),  whereupon  it  is  conducted  through 


FIG.  49. — PREPARATION  OF  HYDROFLUOSILICIC  ACID. 

a  doubly-bent  glass  tube  into  a  cylindrical  jar  containing  a  little  mer- 
cury (into  which  the  tube  opens)  and  on  top  of  this. some  water.  Every 
bubble  of  gas  that  rises  from  the  mercury  into  the  water  generates  in 
the  latter  a  cloud  of  silicic  acid.  If  the  glass  tube  opened  directly  in 
water,  it  would  soon  become  stopped  up  because  of  this  decomposition. 


193.]  OXYGEN  COMPOUNDS  OF  SILICON.  265 

The  solution  of  hydrofluosilicic  acid  reacts  acid;  it  dissolves 
metals  with  the  evolution  of  hydrogen  and  behaves  in  all  respects 
like  an  acid.  A  hydrate,  H2SiF6 + 2H2O;  is  known  in  the  solid  state. 
It  melts  at  19°,  and  is  obtained  by  leading  silicon  fluoride  into  con- 
centrated hydrofluoric  acid.  Most  of  the  salts  of  hydrofluosilicic 
acid  are  soluble  in  water;  the  potassium  salt  is  difficultly  so,  how- 
ever, and  the  barium  salt  is  insoluble. 

Hydrofluosilicic  acid  is  used  in  hardening  objects  made  of  gyp- 
sum (this  is  due  probably  to  the  formation  of  calcium  fluoride)  and 
also  in  analytical  chemistry. 

Oxygen  Compounds  of  Silicon. 

193.  Only  one  such  compound  is  known:  silicon  dioxide,  or 
silica. 

SILICA,    Si02. 

This  compound  occurs  in  astonishingly  large  quantities  and  in 
a  great  number  of  varieties  in  the  solid  crust  of  the  earth.  It  is 
found  crystallized  as  rock  crystal,  quartz  (when  colored  brown, 
called  smoky  quartz),  amethyst  (the  more  beautiful  sorts  being  used 
for  ornament),  tridymite,  onyx,  cat's-eye,  etc.  Sand  is  largely 
silica;  sandstone  also  belongs  here  and  so  does  jasper  (usually 
colored  red  with  ferric  oxide  and  having  a  conchoidal  fracture). 
Opal  is  an  amorphous  variety,  containing  varying  amounts  of 
water. 

Silica  can  be  prepared  artificially  as  an  amorphous  white  powder 
by  heating  silicic  acid. 

Physical  Properties. — In  the  crystallized  state  silica  is  very  hard 
and  insoluble  in  water  and  has  a  specific  gravity  of  2.6.  It  is  very 
difficultly  fusible;  in  the  oxyhydrogen  flame  it  softens  and  passes 
over  into  a  vitreous  modification.  When  heated  strongly  this  can 
be  drawn  out  into  extremely  fine  threads  that  are  so  tenacious  and 
display  so  regular  a  torsion  that  they  are  frequently  used  in  sus- 
pending magnets,  etc.,  in  physical  instruments.  It  can  be  made 
to  boil  vigorously  in  the  electric  furnace;  the  vapor  condenses 
in  woolly  flakes.  Quartz  that  has  been  fused  has  a  very  small 
coefficient  of  expansion  (about  Vi7  of  that  of  platinum);  this 
explains  why  objects  made  of  it  can  endure  very  sudden  changes 
of  temperature.  They  can  be  heated  very  hot  and  then  thrust 


266  INORGANIC  CHEMISTRY.  [§§  193.. 

into  cold  water  at  once  without  cracking.  They  are  attacked  only 
by  metallic  oxides  and  at  a  high  temperature.  Recently  it  has 
become  possible  to  utilize  fused  (vitreous)  quartz  for  the  manu- 
facture of  chemical  apparatus.  It  is  interesting  that  quartz 
vessels  are  transparent  to  ultraviolet  rays,  which  is  not  the  case 
with  glass  vessels. 

Chemical  Properties. — Especially  in  the  crystallized  condition 
silica  is  very  little  acted  upon  by  acids  except  hydrofluoric  acid 
(§  193).  Fused  alkalies  dissolve  it,  forming  alkali  silicates.  It  can 
be  reduced  by  carbon  in  the  electric  furnace,  carborundum  (§  178) 
being  formed.  It  is  also  reduced  by  heating  with  magnesium 
(§  190). 

Silicic  Acids. 

194.  When  a  solution  of  potassium  or  sodium  silicate  (water' 
glass]  is  treated  with  hydrochloric  acid,  a  very  voluminous,  gelat- 
inous mass  separates  out;  this  consists  of  hydrous  silicic  acid  cor- 
responding to  the  general  formula  SiC^aq.  After  being  washed 
with  water  and  dried  in  the  air  it  is  a  fine  white  amorphous  powder 
of  the  approximate  composition  H2Si03.  Freshly  precipitated 
silicic  acid  is  slightly  soluble  in  water,  but  more  so  in  dilute  hydro- 
chloric acid  If,  therefore,  water-glass  is  introduced  into  an  excess  of 
hydrochloric  acid,  the  silicic  acid  stays  in  solution;  it  can  be  sepa- 
rated from  the  sodium  chloride  simultaneously  formed,  by  the  fol- 
lowing process : 

The  solution  is  put  into  a  piece  of  parchment  tubing,  which 
is  tied  at  both  ends,  and  the  whole  submerged  in  pure  water,  the 
latter  being  frequently  renewed.  It  is  found  that  the  salt  goes 
through  the  parchment,  but  that  the  silicic  acid  does  not.  This 
process  is  called  dialysis  and  any  arrangement  for  carrying  it  out  is 
known  as  a  dialyzer.  GRAHAM  found  that  crystallizable  substances 
in  solution  (crystalloids)  are  able  to  pass  through  such  a  membrane, 
while  other  substances,  which  he  called  colloids,  are  not.  In  the 
latter  class  are  glue,  gums,  gelatine,  albumen — in  short,  many 
amorphous  substances  occurring  in  the  animal  and  vegetable 
kingdoms. 

The  silicic  acid  which  separates  from  the  colloidal  solution 
dries  in  the  air  to  a  white  amorphous  powder,  still  containing  a 


195.]  SILICIC   ACIDS.  267 

good  deal  of  water,  however.  The  water  can  be  slowly  extracted 
in  a  sulphuric  acid  desiccator. 

Since  silicon  tetrachloride  is  changed  to  silicic  acid  by  water, 
just  like  phosphorus  pentachloride  to  phosphoric  acid  (§  145),  we 
can  consider  the  compound  as  the  basis  from  which  the  remaining 
silicic  acids  are  derived.  The  latter  can  in  general  be  represented 
by  the  formula  mSi(OH)4— nH20. 

These  polysilicic  acids  themselves  have  not  been  isolated,  but 
many  of  their  salts  and  double-salts  are  known,  which  occur  as 
minerals  in  nature. 

The  silicates  of  potassium  and  sodium  are  soluble  in  water,  those 
of  the  other  metals  insoluble,  as  are  also  most  of  the  double  silicates 
of  the  alkalies. 

In  the  soil  hydrous  silicates  are  found  whose  bases  are  usually  lime 
and  alumina.  In  contact  with  alkali  salts  these  undergo  a  double  decom- 
position, an  insoluble  potassium  aluminium  silicate,  for  example,  being 
formed  together  with  chloride  of  calcium,  which  is  taken  off  by  the  under- 
ground water.  This  phenomenon  is  said  to  be  caused  by  the  absorptive 
power  of  the  soil;  it  plays  an  important  role  in  the  determination  of  soil- 
values.  It  is  this  that  holds  back  the  potash,  an  invaluable  nutrient, 
which  is  furnished  to  the  soil  in  the  form  of  potassium  salts  and  would 
otherwise  be  quickly  washed  off  by  the  rain  because  of  its  solubility. 

The  soluble  phosphates  are  "absorbed*'  by  the  soil  in  the  same 
way.  This  is  mainly  to  be  ascribed  to  the  lime  they  contain,  with  which 
insoluble  tri-  or  dicalcium  phosphate  is  formed;  to  some  extent  this  ab- 
sorption may  be  caused  also  by  basic  lime  silicates. 


Silicon  Compounds  of  Other  Elements. 

195.  Silicon  sulphide,  SiS2,  is  produced  when  carbon  disulphide  vapor 
is  led  over  a  mixture  of  charcoal  and  silica  at  red  heat.  It  forms  long, 
silken  needles,  which  are  broken  up  by  water  into  SiO2  aq  and  hydrogen 
sulphide.  ; 

Silicon  nitride,  Si2N3,  a  white  amorphous  substance,  results  from  the 
heating  of  silicon  in  an  atmosphere  of  nitrogen.  (For  metal  silicides  cf. 
§  190.) 


268  INORGANIC  CHEMISTRY.  [§ 


COLLOIDS.     • 

196.  In  silicic  acid  we  have  become  acquainted  with  a  sub- 
stance that  occurs  in  a  special  form,  viz.,  as  a  colloid.  A  con- 
siderable number  of  such  substances  is  now  known,  and  the 
study  of  them  has  latterly  been  so  active  and  prolific  that  a 
brief  recapitulation  of  the  principal  results  is  fitting  at  this  point. 

GRAHAM  discovered  that  in  aqueous  solution  a  number  of 
substances,  principally  amorphous  materials,  such  as  the  glues, 
albumen  and  dextrin,  have  a  very  small  power  of  diffusion, 
quite  contrary  to  most  salts.  Accordingly  he  distinguished 
between  colloids  and  crystalloids.  Subsequent  investigations 
served  to  increase  greatly  the  number  of  colloids,  i.e.,  substances 
of  small  diffusibility.  Gradually  the  view  developed  that  the 
colloidal  condition  is  not  something  peculiar  to  certain  com- 
pounds, but  that  all  sorts  of  substances,  even  the  crystalloids, 
can  be  obtained  colloidal  by  suitable  treatment.  Hence  the 
colloidal  state .  is  now  regarded  as  a  general  property  of  matter. 
Just  as  we  have  substances  in  the  solid,  liquid  and  gaseous  states, 
so  we  can  also  transform  them  into  the  colloidal  state. 

The  question  that  arises  first  is:  How  may  this  condition 
be  brought  about?  The  following  methods  serve  the  purpose: 

In  the  first  place,  colloids  may  be  prepared  by  simply  dis- 
solving certain  substances,  such  as  glue,  in  water. 

Secondly,  they  are  formed  in  many  cases  instead  of  pre- 
cipitates, when  no  ions  are  present.  For  example,  if  hydrogen 
sulphide  is  passed  into  a  solution  of  arsenic  trioxide,  there  results, 
instead  of  a  precipitate  of  As2Ss,  a  yellow  liquid  containing  the 
arsenic  sulphide  in  colloidal  solution.  However,  if  the  arsenic 
trioxide  solution  is  first  acidified  with  a  little  hydrochloric  acid 
(a  highly  ionized  substance),  the  As2Ss  separates  out  as  a  yellow 
precipitate.  Again,  we  may  take  mercuric  cyanide,  a  compound 
that  is  hardly  ionized  at  all  in  aqueous  solution.  If  a  solution 
of  it  is  treated  with  hydrogen  sulphide,  which  is  also  a  very  feebly 
ionized  substance,  the  mercuric  sulphide  that  is  formed  is 
retained  in  colloidal  solution;  yet,  the  usual  precipitate  can  be 
obtained  by  adding  previously  a  small  amount  of  a  strong  mineral 
acid. 


§  190.]  COLLOIDS.  269 

A  third  way  of  preparing  colloids  is  by  dialysis,  a  process 
described  in  connection  with  silicic  acid.  In  this  way  hydrosols 
(see  below)  of  ferric  oxide,  aluminium  oxide  and  many  other 
substances  can  be  obtained.  Ferric  oxide  hydrosol,  for  instance, 
is  formed  when  ferric  chloride,  FeCls,  is  dissolved  in  water,  and 
just  a  little  less  ammonia  added  than  would  produce  a  precipitate, 
and  the  whole  then  dialyzed.  The  ammonium  chloride,  NH4C1, 
and  hydrochloric  acid  (resulting  from  a  partial  hydrolysis  of 
FeCls  in  aqueous  solution)  pass  through  the  membrane,  while 
Fe^Os  aq.  remains  inside  in  colloidal  solution. 

A  fourth  method  is  the  comminution,  or  dusting,  of  metals 
under  water.  This  is  accomplished  by  connecting  wires  or  rods 
of  platinum,  gold  and  other  metals  with  the  poles  of  a  110-volt 
circuit;  if  the  wires  are  moved  toward  each  other  under  water, 
a  small  arc  is  formed  when  they  are  a  short  distance  apart,  and 
dark  clouds  of  the  metal  proceed  out  into  the  liquid  from  the 
cathode.  The  liquid  is  then  filtered;  the  coarser  bits  of  metal 
remain  on  the  filter  and  the  filtrate  is  a  clear,  dark-colored  solu- 
tion containing  the  metal  as  hydrosol. 

Metals  can  often  be  converted  into  the  colloidal  state  by  treat- 
ing a  very  dilute  solution  of  one  of  their  salts  with  certain  reduc- 
ing-agents  at  ordinary  temperature.  Thus  from  a  very  dilute 
gold  chloride  solution  the  colloidal  gold  can  be  prepared  by 
the  addition  of  phenylhydrazine  hydrochloride  or  acetylene. 

Finally,  it  is  worthy  of  note  that  by  means  of  protective  colloids 
many  substances  can  be  obtained  colloidal  when  other  means 
fail.  For  instance,  if  a  silver  nitrate  solution  and  a  potassium 
bromide  solution,  each  containing  about  1%  of  gelatine,  are 
mixed  together,  the  silver  bromide  is  not  precipitated,  but  comes 
out  colloidal.  "  Collargol,"  a  therapeutic  preparation,  is  a  silver 
colloid,  made  stable  by  a  protective  colloid. 

Colloids  can  be  divided  into  two  groups,  reversible  and 
irreversible.  The  reversible  colloids  comprise-  among  other  sub- 
stances the  agglutinants ,  as  they  are  called, — gelatine,  agar-agar, 
albumins,  starch,  etc.  When  they  are  mixed  with  water  they 
swell  up  and  on  being  gently  warmed  form  a  solution.  When 
cooled  they  gelatinize,  i.e.,  they  congeal  to  a  soft,  viscous  mass 
which  retains  all  the  solvent  water.  The  solution  itself  is  called 


270  INORGANIC  CHEMISTRY.  [§  196. 

a  hydrosol  (or,  in  case  alcohol  is  the  solvent,  an  alcosol)  and  the 
gelatinized  mass  a  hydro  gel. 

When  the  solvent  water  is  extracted  from  a  reversible  colloid 
by  evaporation  at  a  low  temperature  a  hydrogel  is  at  first  formed, 
which  still  contains  a  great  deal  of  water.  This  water  is  partially 
lost  on  exposure  to  the  air.  More  rapidly  in  a  desiccator, — and 
its  vapor  tension  does  not  differ  perceptibly  from  that  of  pure 
water.  When,  however,  a  certain  stage  of  dehydration  is  reached 
the  vapor  tension  begins  to  dimmish.  If  water  is  added  to  the 
hydrogel  before  this  stage  is  reached,  a  hydrogel  is  again  obtained 
with  the  same  properties  as  originally.  The  process  of  s  o  1- 
and  g  e  1-formation  is  thus  a  reversible  one. 

Other  interesting  properties  are  attached  to  hydrosols. 
Crystalloid  salts,  for  example,  diffuse  in  them, — even  in  the  con- 
gealed mass, — almost  as  easily  as  in  water.  If  a  piece  of  jellied 
agar-agar  is  immersed  for  some  time  in  a  dark  blue  ammoniacal 
solution  of  a  copper  salt,  the  agar-agar  becomes  stained  through- 
out its  entire  mass.  Colloids,  on  the  contrary,  do  not  diffuse. 
This  can  be  shown  by  a  colloidal  solution  of  Prussian  blue,  which 
does  not  penetrate  at  all  into  the  agar-agar,  as  above.  The 
electrical  conductance,  too,  is  practically  the  same  for  a  gel 
containing  crystalloid  salts  in  solution  as  for  an  aqueous  solu- 
tion of  the  same  salts  at  like  concentration. 

Oftentimes  large  amounts  of  crystalloid  salts  can  be  added  to 
a  reversible  hydrosol  without  the  formation  of  gel.  This  is  very 
different  with  the  second  class  of  colloids,  the  irreversible  colloids, 
for  they  are  in  many  cases  very  sensitive  to  additions  of  salts. 
When  the  salt  is  added  the  irreversible  hydrosol  begins  to  appear 
cloudy,  and  a  precipitate  is  formed  which  cannot  be  reconverted 
offhand  into  hydrosol. 

Irreversible  hydrosols  can  be  prepared  in  the  various  ways 
already  mentioned.  They  comprise  the  colloidal  metals,  sul- 
phides, hydrated  oxides,  etc.  Most  of  them  are  mobile  liquids, 
in  contrast  to  many  reversible  colloids,  such  as  glue. 

The  quantity  of  an  electrolyte  that  is  just  sufficient  to  pre- 
cipitate an  irreversible  hydrosol  is  connected  with  the  valence 
of  the  electrolyte,  the  quantity  decreasing  rapidly  with  in- 
creasing valence.  The  As2S3  hydrosol  is  just  coagulated  by 


§  196.]  COLLOIDS.  271 

71  mill-equivalents  of  NaCl  per  liter,  2.0  of  MgCl2  and  0.39 
of  A1C13. 

Certain  irreversible  colloids  are  capable  of  mutually  pre- 
cipitating each  other;  others  are  not.  The  hydrosols  of  ferric 
oxide  and  arsenious  sulphide  give  a  precipitate  when  mixed, 
but  the  hydrosols  of  gold  and  arsenious  sulphide  mix  without 
precipitation.  These  phenomena  have  been  shown  to  be  connected 
with  the  behavior  of  the  substances  toward  the  electric  current. 
If  a  solution  is  introduced  into  a  U-tube  supplied  with  electrodes 
at  the  upper  ends  and  a  strong  current  (say  110  volts)  is  passed 
through  it,  the  colloid  is  seen  to  separate  out  and  wander  either 
to  the  anode  or  to  the  cathode.  At  one  of  the  two  electrodes  an 
aqueous  layer  appears,  which  is  entirely  free  from  colloid  and  is 
separated  sharply  from  the  hydrosol.  This  convective  trans- 
ference, or  electrical  endosmose,  is  by  no  means  to  be  confused 
with  the  ionic  migration  in  electrolytes.  For,  while  in  the  elec- 
trolytes there  is  an  electrical  opposition  between  the  dissocia- 
tion products  of  the  dissolved  substance,  the  electrical  opposi- 
tion exists  in  this  case  between  the  colloid  and  the  solvent. 
In  general,  mutual  precipitation  is  only  possible  with  colloids 
whose  electrical  charges  are  opposite  with  respect  to  that  of  a 
common  solvent. 

The  colloidal  state  must  be  regarded  as  a  very  fine  divi- 
sion, or  distribution,  of  one  substance  in  another.  This  follows 
from  the  great  analogy  which  exists  between  suspensions  (e.g., 
clay  and  water)  and  colloids.  Both  can  be  separated  out 
by  centrifuging  and  both  display  the  TYNDALL  effect.  This 
effect  may  be  described  as  follows:  When  a  beam  of  light 
passes  through  a  body  of  air  that  is  free  from  dust  it  is 
invisible  transversely;  the  gas  is  "  optically  a  vacuum/'  But, 
so  soon  as  dust  particles  enter  the  air,  the  path  of  the  beam  can 
be  followed  through  the  dispersion  of  the  light  by  the  particles. 
Optically  vacuous  liquids  and  optically  vacuous  solutions  of 
crystalloid  salts  can  also  be  prepared.  But  if  a  beam  of  light  is 
passed  through  a  hydrosol  the  path  of  the  beam  can  be  seen. 
The  hydrosol  is  therefore  not  an  optical  vacuum;  it  must  con- 
tain floating  particles,  but  these  are  so  small  that  they  can  not  be 
seen  even  with  the  best  microscopes. 


272  INORGANIC  CHEMISTRY.  [§§  196- 

However,  SIEDENTOPF  and  ZSIGMOXDY  have  succeeded  in 
rendering  these  sub  microscopic  particles  visible  with  an 
apparatus  that  they  call  the  ultramicroscope.  In  it  the  hydrosol 
is  illuminated  transversely  so  that  the  luminous  rays  do  not  blind 
the  eye  of  the  observer.  The  submicroscopic  particles  bend  (dif- 
fract) the  light  rays "  in  all  directions,  so  that  with  sufficiently 
intense  illumination  the  light  effect  produced  by  each  individual 
particle  comes  within  the  range  of  microscopic  visibility  and  can 
be  separately  observed  without  however  revealing  its  form. 

Suspensions  resemble  colloids  further  in  that  they  exhibit 
electrical  endosmose  and  can  be  precipitated  by  the  addition  of 
electrolytes. 

When  a  liquid  is  distributed  through  another  in  exceedingly 
gmall  drops  we  have  an  emulsion.  The  most  familiar  example  is 
milk,  an  emulsion  of  butter  fat.  There  is  reason  for  assuming 
that  many  reversible  colloids  are  extremely  fine  emulsions;  for 
instance,  an  emulsion,  like  a  gelatine  solution,  cannot  be  coag- 
ulated by  the  addition  of  an  electrolyte. 

The  knowledge  that  in  the  colloidal  state  we  have  to  do  with 
a  very  fine  distribution  of  one  substance  in  another  has  led  to  the 
introduction  of  a  new  set  of  terms.  The  substance  distributed 
as  a  colloid  is  now  generally  spoken  of  as  the  disperse  phase, 
distributed  in  the  dispersion  medium.  Further,  the  words  dis- 
persoids  and  emulsoids  are  replacing  the  word  "colloids." 

It  was  formerly  thought  that  a  sharp  distinction  must  be 
drawn  between  colloids  and  real  solutions.  GRAHAM  spoke  of 
two  different  worlds  of  matter.  In  contrast  to  the  true  solu- 
tions colloids  exhibit  practically  no  diffusion,  no  vapor  pressure 
lowering,  no  boiling-point  elevation  or  freezing-point  depres- 
sion,— in  short,  no  osmotic  phenomena.  The  researches  of  recent 
years  have,  however,  shown  that  essential  differences  do  not 
really  exist.  To  begin  with,  we  have  come  upon  many  cases 
of  transition  between  colloidal  and  real  solutions.  Furthermore, 
it  was  previously  observed  by  LOBRY  DE  BRUYN  that  salt  solu- 
tions can  be  separated  by  centrifugal  force  into  portions  of  unlike 
concentration.  More  important  still,  the  investigations  of 
EINSTEIN,  PERRIN,  SVEDBERG  and  others  have  shown  that 
colloids,  just  like  true  solutions,  are  subject  to  the  osmotic  laws. 


197.]  GERMANIUM.  273 

From  the  molecular-kinetic  point  of  view  there  is  no  difference, 
according  to  these  investigations,  between  a  "  dissolved  molecule  ' 
and  a  "  suspended  particle";  consequently  a  mechanical  sus- 
pension must  exert  exactly  the  same  osmotic  pressure  as  a  "  true 
solution  "  of  the  same  number  of  particles  per  unit  volume. 
The  fact  that  the  colloidal  solutions  display  no  properties  cor- 
responding to  osmotic  pressure  is  simply  due  to  the  fact  that  at 
the  same  concentration  the  number  of  freely  moving  particles 
in  solutions  is  enormously  greater  than  with  the  colloids,  i.e.,  an 
individual  colloid  particle  has  gigantic  dimensions  as  compared 
with  those  of  a  molecule.  This  makes  the  freezing-point  lower- 
ing, etc.,  so  slight  that  it  cannot  be  measured  by  present  exper- 
imental means. 

In  order  to  test  the  applicability  of  the  osmotic  laws  to  colloidal 
solutions  we  are  therefore  forced  to  employ  indirect  methods. 
The  methods  employed  are  associated  with  four  phenomena: 

(1)  the  translatory  and   rotatory   movements  of  the  particles. 

(2)  diffusion ;   (3)   the  change  of  concentration  under  the  influence 
of  gravity;  and  (4)  the  local  temporary  changes  of  concentration. 
The  possibility  of  testing  the  osmotic  laws  by  such  measurements 
is  a  result  of  developing  formulae  for  these  phenomena  that  are 
deduced  upon  the  assumption  that  the  osmotic  laws  are  applic- 
able. 

These  investigations  also  serve  to  corroborate  the  reality  of 
atoms  and  molecules,  supporting  the  information  gained  in  many 
other  ways,  as  has  been  set  forth  in  §  35. 

GERMANIUM. 

197.  This  element  is  of  extremely  rare  occurrence.  It  was  discovered 
by  WINKLER  in  an  argentiferous  mineral,  argyrodite,  GeS2-4Ag2S,  found 
in  Freiberg,  in  Saxony.  Germanium  forms  grayish-white  octahedrons  with 
a  metallic  lustre  and  a  specific  gravity  of  5.469  at  20°.  It  melts  at  900°. 
At  ordinary  temperatures  it  is  unaffected  by  the  air;  at  red  heat  it  burns, 
forming  white  fumes  of  germanium  oxide,  GeO2.  Two  series  of  compounds 
of  this  element  are  known,  which  are  derived  from  the  oxides  GeO  and 
GeO2;  the  ous  compounds  are  easily  oxidized  to  the  higher  form,  germanic 
acid.  The  hydrogen  compounds,  GeH4  and  GeHCl3,  are  known. 

Germanic  chloride,  GeCl4,  can  be  prepared  directly  from  the  elements. 
It  is  broken  up  by  water  forming  Ge(OH)4. 


274  INORGANIC  CHEMISTRY.  [§§  197- 

Germanium  dioxide,  GeO2,  is  produced  by  heating  the  corresponding 
hydroxide,  or  by  roasting  the  element  or  its  sulphide  or  by  treating  it  with 
nitric  acid.  It  is  a  white  powder  of  a  specific  gravity  of  4.703  at  18°  and 
is  unaffected  by  heat. 

Germanium  disulphide,  GeS2,  separates  as  a  white  precipitate  when 
hydrogen  sulphide  is  passed  into  the  solution  of  germanium  dioxide  in 
strong  hydrochloric  acid.  In  moist  air  it  decomposes,  giving  off  hydrogen 
sulphide.  It  dissolves  in  alkalies  and  alkali  sulphides  to  form  sulpho-salts. 

For  germanium  cf.  also  §  218. 

TIN. 

198.  This  metal  is  not  very  widely  distributed  on  the  earth;  in 
some  places,  however,  it  is  found  in  quite  large  quantities.  The 
principal  tin  mines  of  Europe  are  those  in  Cornwall;  even  the 
Phoenicians  obtained  tin  there.  The  most  important  present  locali- 
ties are  on  the  group  of  islands  lying  east  of  Sumatra  (Banca,  Bil- 
liton,  Sinkop,  etc.).  There  the  metal  occurs  in  the  form  of  tin- 
stone (cassiterite,  SnC^);  it  is  found  in  quadratic  crystals,  which 
are  usually  colored  brown  or  black  by  a  small  amount  of  iron.  In 
order  to  extract  the  metal,  the  ore  is  at  first  roasted,  to  eliminate 
any  sulphur  or  arsenic  it  may  contain,  and  then  reduced  with  car- 
bon. The  tin  thus  obtained  is  refined  by  liquation,  i.e.  by  fusing 
again  at  a  low  temperature  and  pouring  it  off  from  the  less  fusible 
alloy  of  tin  with  iron  and  arsenic.  It  is  then  melted  once  more 
and  stirred  with  a  wooden  pole  (branch  of  a  tree),  whereby  the 
oxide  still  remaining  is  reduced.  The  Banca  tin  is  nearly  chemic- 
ally pure. 

Physical  Properties. — Tin  is  a  silvery-white  metal,  melting  at 
232.7°  and  volatilizing  between  1450°  and  1600°.  Sp.  g.  =  7.293 
at  13°.  It  has  a  crystalline  structure  which  can  be  made  visible 
by  moistening  with  hydrochloric  acid,  whereupon  peculiar  frost- 
like  etch-figures  are  produced  on  the  surface  (tin-moiree).  When  tin 
is  bent,  a  characteristic  crackling  sound  (cry  of  tin)  is  heard,  which 
is  probably  caused  by  the  grating  of  the  crystal  faces  on  each  other. 
Tin  is  very  malleable  and  ductile;  it  can  be  beaten  into  very  thin 
leaves  (tin-foil)  at  the  ordinary  temperature,  and  at  100°  it  can  be 
drawn  out  into  wire.  At  a  very  low  temperature  and  in  contact 
with  an  alcoholic  pink-salt  solution  (§  201),  tin  passes  spontaneously 
into  another  modification,  gray  tin,  which  has  a  lower  specific  gravity, 


198.J  TIN.  275 

5.8.  Above  20°  this  form  changes  back  to  white  tin.  If  the  latter 
is  brought  in  contact  with  gray  tin  at  ordinary  temperatures 
(below +  20°),  it  turns  very  slowly  into  gray  tin,  falling  to  powder, 
probably  because  of  the  increase  in  volume  (this  phenomenon  is 
called  the  "tin-disease")-  If  it  is  not  in  contact  with  the  gray 
modification,  the  transformation  does  not  take  place  at  all  at 
ordinary  temperatures,  or  at  least  not  for  centuries.  Evidently 
there  is  a  transition  point  of  the  two  forms  at  20°,  and  we  are 
forced  to  the  odd  conclusion  that,  except  on  warm  summer  days, 
tin  is  in  the  metastable  condition. 

The  reason  why  tin,  even  in  contact  with  gray  modification, 
passes  so  slowly  into  that  form  at  ordinary  temperatures  is  that 
the  velocity  of  transformation  is  small  in  the  neighborhood  of  the 
transition  point;  it  is  accelerated  on  moving  away  from  that  point. 
When  the  temperature  sinks  this  acceleration  is  counteracted, 
however,  by  the  retardation  that  all  reactions  undergo  by  a  lower- 
ing of  temperature.  In  many  cases,  therefore,  there  must  be  a 
maximum  of  the  velocity  of  transformation,  such  as  we  have  here 
at  —48°;  below  that  temperature  the  transformation  again  becomes 
slower. 

Ordinary  tin  crystallizes  in  the  tetragonal  system.  In  addition  to  tlie 
gray  modification  there  is  also  a  third  one,  the  rhombic  modification.  The 
transition  point  tetragonal  ^rhombic  is  about  170°. 

This  point  was  determined  in  a  unique  way,  namely,  by  measuring  the 
velocity  of  flow  of  the  metal  under  high  pressure.  For  this  purpose  the 
solid  metal  was  placed  in  a  cylinder  having  a  hole  in  the  bottom,  and  the 
quantity  of  metal  was  measured  that  was  forced  out  under  constant  pres- 
sure in  the  unit  of  time.  In  general,  this  quantity  increases  rapidly  with 
rising  temperature,  but  with  tin  it  was  found  to  diminish  considerably  when 
the  temperature  reached  about  200°.  This  may  be  taken  as  a  proof  that 
the  metal  has  another  (third)  modification.  At  200°  tin  is  so  brittle  that 
it  can  be  easily  pulverized. 

Chemical  Properties. — Tin  is  unaffected  by  the  air  at  ordinary 
temperatures;  if  heated  strongly,  it  burns  with  an  intense  white 
light  to  tin  oxide,  Sn02.  Hydrochloric  acid  dissolves  it,  forming 
stannous  chloride  and  hydrogen.  It  is  also  attacked  by  nitric  acid 
(§  201).  A  boiling  solution  of  caustic  soda  or  potash  converts  it 
into  a  stannic  acid  salt  (s  t  a  n  n  a  t  e)  with  the  evolution  of  hydro- 
gen: 

Sn + 2KOH  +  H2O  =  K2SnO3  +  2H2. 


276  INORGANIC  CHEMISTRY.  [§§  199- 

In  the  presence  of  weak  acids  (acetic  acid)  and  alkalies  it  is  very 
stable. 

199.  Uses.  —  On  account  of  its  permanence  tin  is  used  as  a  pro- 
tective covering  for  metals  which  are  attacked  by  the  air  and  the 
above-named  agencies.    Many  kitchen  utensils  are  "  tinned."     Sheet 
iron  is  covered  with  a  layer  of  tin,  to  protect  it  from  rusting  (§  279), 
and  is  then  known  as  tin-plate,  or  sheet-tin.    This  is  done  by  simply 
dipping  the  sheet  iron,  which  has  been  cleaned  by  hydrochloric 
or  sulphuric  acid,  in  molten  tin. 

Many  alloys  of  tin  are  in  use.  Solder  consists  of  tin  and  lead 
(in  the  ratio  2:1  or  1:1  or  1:2),  and  is  harder  than  either  of  its 
components  but  more  easily  fusible.  The  alloys  of  c  o  p  p  e  r  and 
t  i  n  are  called  bronzes;  their  composition  varies  according  to  the 
purpose  they  serve.  At  present  the  bronzes  usually  contain  a 
little  lead  and  zinc  as  well.  Bronze  is  hard  and  tough,  can  be 
easily  worked  and  fuses  to  a  mobile  liquid,  hence  it  is  particularly 
suitable  for  casting.  Gun  metal  contains  90%  copper  and  10%  tin; 
bell  metal  20-25%  tin,  the  rest  being  copper.  Phosphor  bronze  is 
prepared  by  fusing  copper  with  tin  phosphide  (§  202).  The  result- 
ing mass  is  remarkably  homogeneous  and  contains  0.25-2.5% 
phosphorus  and  5-15%  tin.  Its  great  hardness  and  firmness  render 
it  especially  valuable  for  certain  parts  of  machines  (axle-bearings). 
Silicon  bronze  contains  silicon  in  place  of  phosphorus,  is  very  hard 
and  conducts  electricity  well,  hence  it  is  used  for  making  telephone 
wire.  Tin  amalgam  forms  the  metallic  coating  of  mirrors. 

Compounds  of  Tin. 

Tin  forms  two  sets  of  compounds;  they  correspond  to  the 
oxygen  compounds,  stannous  oxide,  SnO,  and  stannic  oxide,  Sn02. 

STANNOUS  COMPOUNDS. 

200.  Stannous  chloride,  SnCl2,  is  prepared  by  dissolving  tin  in 
hydrochloric  acid: 


It  crystallizes  with  two  molecules  of  water,  which  are  given  off  at 
100°.  :-  It  is  very  readily  soluble  in  water  (1  part  in  0.37  at  ordinary 
temperatures).  Anhydrous  stannous  chloride  is  white  and  trans- 
parent; it  melts  at  250°  and  boils  at  606°.  A  little  above  the 


200.]  STAXXOUS   COMPOUNDS.      ,  277 

boiling-point  the  vapor  density  corresponds  to  the  formula  Sn2CLi; 
above  900°,  however,  to  SnCl2. 

The  aqueous  solution  acts  strongly  reducing.  It  absorbs  oxy- 
gen from  the  air  with  the  partial  formation  of  basic  chloride  (a 
white  powder),  if  the  liquid  is  not  too  acidic: 

3SnCl2  +  H20  +  0  =  SnCl4  +  2Sn(OH)Cl. 

Basic  chloride. 

But  if  the  liquid  is  strongly  acid  the  tetrachloride  SnCU  is  also 
formed  in  this  oxidation. 

This  same  basic  chloride  also  results  from  hydrolytic  dissociation, 
when  a  neutral  stannous  chloride  solution  is  strongly  diluted. 

SnCl2  +aq  =Sn(OH)Cl  +  HC1  +aq. 

The  reducing  power  of  stannous  chloride  is  further  seen  in  its 
action  on  potassium  permanganate,  potassium  dichromate,  cupric 
chloride,  mercuric  chloride,  etc.,  all  of  which  are  converted  into 
lower  stages  of  oxidation  in  acid  solution. 

It  may  be  remarked  here  that,  from  the  ionic  point  of  view, 
oxidation  amounts  in  many  cases  to  raising  an  ion  to  a  higher 
positive  potential,  and  reduction  to  the  reverse.  Let  us  consider, 
for  instance,  the  reaction  between  stannous  chloride  and  mercuric 
chloride.  This  can  be  expressed  by  the  equation 

SnCl2  +  HgCl2  =  SnCl4  +  Hg. 

Stannous  chloride  is  oxidized  to  stannic  chloride;  at  the  same 
time  mercuric  chloride  is  "  reduced  "  to  the  metal.  Written  in 
ions,  this  equation  becomes 


that  is,  the  electrical  charge  of  the  mercury  ion  is  taken  by  the 
bivalent  tin  ion,  the  former  losing  its  electrification. 

Another  example  is  the  action  of  chlorine  on  stannous  chloride, 
by  which  the  latter  is  "oxidized"  to  stannic  chloride: 


=  SnCl4. 
The  ionic  reaction  is 


278  INORGANIC   CHEMISTRY.  [§§  200- 

Tin  takes  up  two  more  positive  charges,  but  this  necessitates  that 
the  two  Cl  atoms  become  ions;  they  thus  require  two  negative 
charges;  but  when  these  are  formed  two  positive  charges  are  ob- 
tained at  the  same  time.  However,  the  Sn*"'  and  Cl'  ions  unite 
to  form  stannic  chloride,  SnCU,  which  is  a  very  weak  electrolyte 
(cf.  §  201). 

In  the  preparation  of  chlorine,  hydrochloric  acid  is  "  oxidized  " 
by  manganese  dioxide  : 

Mn02  +  4HC1  =  MnCl2  +  2H20  +  C12, 
or 


the  positive  charge  of  the  four  H'  ions  is  thus  transferred,  half  to 
the  manganese  and  the  rest  serving  to  discharge  two  chlorine  ions, 
i.e.  to  equalize  their  negative  charges. 

Various  double  salts  of  stannous  chloride  are  known,  e.g. 
SnCl2  •  2KC1  ;  SnCl2  •  2NH4C1. 

Stannous  hydroxide,  Sn(OH)2  is  precipitated  when  a  solution 
of  stannous  chloride  is  treated  with  soda  : 

SnCl2+Na2C03+H2O  =  Sn(OH)2+2NaCl+C02. 

This  hydroxide  is  insoluble  in  ammonia,  but  soluble  in  alkalies; 
when  the  latter  solution  is  boiled,  tin  is  deposited  and  alkali  stan- 
nate,  e.g.  K2SnO3,  formed.  The  hydroxide  is  also  soluble  in 
acids,  thus  displaying  a  basic  as  well  as  an  acidic  nature.  Such 
compounds  are  able  to  give  hydroxyl  ions  (Sn"+2OH/)  on  the 
one  hand  and  hydrogen  ions  (Sn02"  +  2H')  on  the  other.  They 
are  termed  amphoteric  compounds. 

Stannous  oxide  is  obtained  by  heating  the  hydroxide  in  a  cur- 
rent of  carbon  dioxide;  it  is  a  dark-brown  powder,  which  takes 
fire  in  the  air,  burning  to  stannic  oxide,  Sn02. 

Other  salts  of  stannous  oxide  than  the  above-mentioned  stannous 
chloride  are  also  known.  The  sulphate,  for  instance,  is  obtained  by  dis- 
solving the  hydroxide  or  the  metal  in  dilute  sulphuric  acid.  It  forms  a 
basic  salt  readily. 

Stannous  sulphide,  SnS,  is  precipitated  as  an  amorphous  brown 
powder  when  hydrogen  sulphide  is  passed  into  the  solution  of  stan- 
nous salts.  It  is  insoluble  in  potassium  sulphide,  K2S,  but  it 


201.]  STANNIC  COMPOUNDS.  279 

dissolves  to  form  a  sulpho-stannate  when  brought  in  contact 
with  the  p  o  1  y  sulphide  of  ammonium  or  potassium,  K2Sx(x=2-5). 

SnS  +  K2S2  =  K2SnS3. 

Stannous  sulphide  can  also  be  prepared  by  fusing  tin  with  sul- 
phur. It  then  forms  a  bluish-gray  crystalline  mass. 

STANNIC  COMPOUNDS. 

201.  Stannic  chloride,  SnCl4,  was  prepared  as  early  as  1605.  It 
was  named  spiritus  jumans  Libami,  after  its  discoverer.  It  is 
obtained  by  the  action  of  chlorine  on  tin  or  stannous  chloride. 
Stannic  chloride  is  a  liquid  which  fumes  strongly  in  the  air;  it 
boils  at  113.9°,  and  has  a  specific  gravity  of  2.234  at  15°.  When 
brought  in  contact  with  a  little  water  or  on  taking  up  moisture 
from  the  air,  it  goes  over  into  a  semi-solid,  crystallized  mass, 
SnCl4-3H2O,  the  so-called  tin-butter.  A  fresh  solution  of  stannic 
chloride  is  a  very  poor  conductor  of  electricity.  However,  the  con- 
ductivity increases  slowly  at  ordinary,  faster  at  higher,  tempera- 
tures; after  several  days -it  reaches  a  maximum.  In  the  case  of 
more  dilute  solutions  this  maximum  is  much  higher.  These  facts 
can  be  explained  by  assuming  that  stannic  chloride  is  but  feebly 
ionized  and  that  it  reacts  with  water  in  the  following  way; 

SnCl4+4H20  <=>  Sn(OH)4+4HCl; 

in  other  words,  that  it  undergoes  hydrolytic  dissociation.  It  is 
the  liberated  hydrochloric  acid  that  causes  the  conductivity. 
The  solution  contains  tin  hydroxide  in  the  colloidal  state.  The 
water  has  thus  split  up  the  stannic  chloride  into  a  basic  hydroxide 
and  an  acid. 

Stannic  chloride  forms  well-crystallized  double  salts  with  the 
alkali  chlorides,  e.g.  SnCl4-2KCl  and  SnCl4-2NH4Cl.  The  latter 
is  known  as  pink  salt  (because  of  its  color)  and  is  used  as  a  mor- 
dant in  dyeing.  Tin  tetrachloride  also  unites  with  the  chlorides  of 
the  metalloids  to  form  crystallized  substances,  e.g.  SnCU-PCls; 
SnCl4  -  POC13 ;  SnCl4  •  SC14,  etc.  It  combines  with  hydrochloric  acid, 
forming  a  leafy-crystalline  mass,  H2SnCl6  •  6H20,  which  melts  at  9°. 

Tin  fluoride,  SnF4,  itself  is  not  known,  but  there  is  a  compound, 
K2SnF6,  which  corresponds  to  potassium  fluo-silicate;  the  salts  of  hydro- 
fluostannic  acid  are  isomorphous  with  the  analogous  silicon  compounds. 


280  INORGANIC  CHEMISTRY.  [§§201- 

Stannic  oxide,  SnO2,  can  be  prepared  synthetically  by  heating 
tin  in  air.  It  is  an  amorphous  white  powder,  insoluble  in  acids  and  al- 
kalies; the  latter,  however,  dissolve  it  when  f  used,  forayngstannates. 

Stannic  Acid  and  Metastannic  Acid. — The -hydroxides  corre- 
sponding to  SnO2  have  only  very  weakly  basic  properties;  here  the 
acidic  properties  are  prominent.  The  normal  hydroxide,  Sn(OH)  4, 
is  unknown,  but  there  is  a  hydroxide  of  the  empirical  composi- 
tion H2SnO3(— Sn(OH)4—  H2O),  corresponding  to  carbonic  acid, 
H2CO3.  Strangely  enough  this  exists  in  two  modifications,  which 
differ  from  each  other  both  chemically  and  physically;  they  are 
called  stannic  and  metastannic  acids. 

The  stannic  acid  is  precipitated  when  ammonia  is  added  to  an 
aqueous  solution  of  stannic  chloride  or  hydrochloric  acid  to  a  potas- 
sium stannate  solution.  This  precipitate  reacts  acid  when  moist 
and  is  soluble  in  concentrated  hydrochloric  and  nitric  acids,  as  well 
as  in  alkalies.  It  gradually  changes  into  metastannic  acid. 

Metastannic  acid  is  generally  prepared  by  treating  tin  with 
strong  nitric  acid;  it  is  then  formed  in  a  vigorous  reaction  as  a 
dense  white  powder.  Metastannic  acid  is  insoluble  in  sodium 
hydroxide,  but  nevertheless  unites  with  it  to  form  sodium  metastan- 
nate ;  this  is  dissolved  by  water,  although  with  difficulty,  but  is  insolu- 
ble in  the  caustic  soda  solution.  When  boiled  with  hydrochloric 
acid,  metastannic  acid  goes  over  into  a  chloride,  which  is  insoluble 
in  the  concentrated  acid  but  soluble  in  water.  This  solution  does 
not  contain  the  ordinary  tin  chloride,  but  another  one,  meta-tin 
chloride,  having,  however,  the  same  composition,  SnCl4.  It  is 
distinguished  from  the  ordinary  stannic  chloride  by  giving  a  yellow 
coloration  with  stannous  chloride  solution;  the  solution  of  the  ordi- 
nary chloride  does  not  do  this  till  after  some  time,  during  which 
the  metachloride  is  formed  in  it. 

Stannic  acid  and  the  corresponding  chloride  thus  pass  over  into 
the  meta-compounds  spontaneously;  on  the  other  hand,  metastan- 
nic acid  can  be  converted  into  the  ordinary  tin  compounds  by  boil- 
ing it  for  some  time  or  fusing  it  with  a  caustic  alkali. 

The  difference  between  stannic  and  metastannic  acids  was  pointed  out  by 
BERZELIUS  as  early  as  the  beginning  of  the  nineteenth  century.  They  are 
both  colloids.  The  salts  of  metastannic  acid  have  in  general  a  very  com- 
plicated composition,  similar  to  the  polysilicates  (§  195),  for  which  reason 
metastannic  acid  is  regarded  as  a  polymer  of  the  ordinary  stannic  acid,  i.e., 
that  its  molecule  is  represented  by  (H2SnO3)z,  stannic  acid  itself  being  H2SnO3. 


203.]  LEAD.  281 

Of  the  salts  of  stannic  acid,  the  sodium  stannate,  Na2SnO3  + 
3H2O,  is  especially  well  known.  It  comes  on  the  market  under 
the  name  of  "  preparing-salt "  and  is  used  as  a  mordant  in  dyeing. 
It  is  made  by  fusing  tin-stone  with  caustic  soda  and  crystallizes 
in  hexagonal  crystals,  which  are  more  soluble  in  cold  than  in 
warm  water. 

Purple  of  Cassius  is  obtained  when  a  mixture  of  the  hydrosols  of  tin 
dioxide  and  gold  is  precipitated  by  adding  some  such  electrolyte  as  am- 
monium chloride.  This  mode  of  formation  proves  that  the  substance  is 
not  a  compound  of  the  two  components,  as  was  formerly  believed,  but 
only  a  mixed  gel. 

202.  Stannic  sulphide,  SnS2,  falls  out  as  a  yellow  amorphous 
powder,  when  hydrogen  sulphide  is  passed  into  the  acid  solution  of 
a  stannic  compound.     It  can  be  synthesized  by  heating  tin  amalgam 
with  sulphur  and  ammonium  chloride,  being  thus  obtained  in  the 
form  of  transparent  golden  leaves  and  known  as  aurum  musivum, 
or  mosaic  gold;  it  is  used  for  gilding.     Stannic  sulphide  is  a  sulpho- 
anhydride;    the  corresponding  sulpho-acid,  H^SnSs,  is  not  known 
in  the  free  state,  but  exists  in  the  form  of  salts. 

Sodium  sulphostannate,  Na2SnSs  +  2H2O,  crystallizes  in  color- 
less octahedrons.  When  its  solution  is  treated  with  an  acid,  stan- 
nic sulphide  is  precipitated. 

Tin  phosphide  serves,  as  was  stated  above,  for  the  manufacture  of 
phosphor  bronze.  Of  the  various  compounds  of  tin  and  phosphorus, 
the  best  known  is  the  compound  Sn9P.  It  forms  a  coarsely  crystalline 
mass,  which  melts  at  170°. 

LEAD. 

203.  Among  the  lead  ores  the  most  important  is  qalenite  (PbS) ; 
it  occurs  in  isometric  crystals  (cubes)  of  a  graphitic  color.     Other 
ores  are  cerussite  (PbCOs),  crocoite  (PbCrO^,  wulfenite  (PbMoO4), 
etc.     For  the  extraction  of  the  metal  galenite  is  used  almost  exclu- 
sively.    This  is  roasted  to  convert  the  sulphide  partially  into  oxide, 
and  partially  into  sulphate: 

PbS  +  3O  =  PbO  +  SO2;  PbS  +  2O2  =  PbSO4. 

In  roasting  care  is  taken  that  a  considerable  portion  of  the  ore 
remains  as  sulphide.  On  farther  heating,  the  latter  reacts  with  the 
oxygen  compounds  in  the  following  way: 

=  3Pb+SO2;  and  PbSO4+PbS=2Pb+2S02. 


282  INORGANIC  CHEMISTRY.  [§§  203  - 

Physical  Properties.  —  Lead  is  a  soft  ductile  metal  of  a  bluish 
color.  On  exposure  to  the  air  it  loses  its  lustre  rapidly,  becoming 
coated  with  a  very  thin  layer  of  oxide.  It  has  a  specific  gravity  of 
11.254,  melts  at  327°,  and  boils  at  1525°. 

Chemical  Properties.  —  The  thin  coating  formed  by  the  oxide  on 
the  brilliant  surface  of  the  metal  protects  the  lead  from  further 
attack  by  the  air.  If,  however,  it  is  prepared  in  a  very  finely 
divided  state,  e.g.,  by  heating  lead  tartrate  or  citrate  in  the  absence 
of  air,  it  takes  fire  in  the  air  even  at  ordinary  temperatures.  (Other 
metals  can  be  reduced  in  a  similar  way  to  a  fine  state  of  division, 
whereupon  they  ignite  spontaneously  in  the  air.  A  substance 
which  exhibits  this  phenomenon  is  called  a  pyrophorus.) 
When  lead  is  melted,  it  becomes  coated  with  red  oxide  of  lead; 
by  constantly  removing  the  latter,  the  lead  can  be  entirely  oxidized. 
A  compact  mass  is  unaffected  by  sulphuric  or  hydrochloric  acid, 
but,  when  finely  divided,  it  reacts  to  form  the  corresponding  salts. 
Nitric  acid  easily  dissolves  it  to  form  the  nitrate.  Acetic  acid 
and  various  vegetable  acids  attack  it;  since  all  lead  salts  are  very 
piosonous  and  very  serious  effects  result  from  chronic  poisoning 
with  insignificant  but  successive  amounts,  it  is  not  admissible  to 
use  tin  containing  lead  in  tin-plating  vessels  for  use  in  the  kitchen. 

Zinc  and  iron  precipitate  the  metal  from  solutions.  A  piece  of 
zinc  becomes  covered  with  a  dendritic  crystalline  mass  ("  lead- 
tree  ").  This  reaction  can  be  expressed  by: 

Zn+Pb"=Zn" 


i.e.  zinc  is  changed  into  the  ionic  condition,  and  the  lead  ions  are 
discharged.  How  it  comes  about  that  one  metal  thus  assumes  the 
electrical  charge  of  another  may  be  explained  by  a  hypothesis  of 
NERNST.  His  supposition  is  that  every  metal  on  coming  in  con- 
tact with  water  or  a  solution  tends  to  send  positive  ions  into  it. 
This  emission  of  ions  continues  until  the  positive  charge  acquired  by 
the  solution  and  the  negative  charge  created  on  the  metal  balance 
by  their  mutual  attraction  the  tension  (called  the  electrolytic 
solution-tension)  with  which  the  ions  are  driven  into  the  solution. 
This  tension  differs  considerably  for  different  metals;  for  zinc  it 
is  much  greater  than  for  lead.  When,  therefore,  a  strip  of  zinc 
is  dipped  in  a  lead  solution  it  forces  zinc  ions  into  the  solution  and 
the  zinc  thus  becomes  much  more  negatively  charged  than  would 


204.]  OXIDES  OF  LEAD.  283 

a  piece  of  lead  by  the  emission  of  lead  ions.  The  lead  ions  are 
therefore  attracted  by  the  zinc  and  discharged,  i.e.  lead  is  precipi- 
tated from  the  solution.  This  process  stops  only  when  all  the 
lead  of  the  solution  has  been  replaced  by  zinc. 

Distilled  water,  from  which  the  air  has  been  entirely  removed 
by  boiling,  has  no  effect  on  lead,  but  the  simultaneous  action  of  air 
and  water  produce  lead  hydroxide,  which  is  somewhat  soluble  in 
water.  This  hydroxide  is  converted  into  insoluble  basic  carbonate 
by  carbonic  acid. 

From  a  hygienic  standpoint  these  properties  of  lead  are  of  vast  impor- 
tance, because  drinking-water  is  almost  universally  conducted  through 
pipes  made  of  lead  or  material  containing  lead  ("compo-pipes").  The 
absorption  of  lead  from  such  pipes  by  water  and  the  continuation  of 
the  process  depends  in  a  large  measure  on  the  proportion  of  salt  in  the 
water.  As  a  rule,  the  less  of  salts  it  contains,  the  more  lead  it  takes 
up.  Rain-water,  which  is  almost  entirely  free  from  solid  matter,  but 
contains  oxygen,  carbon  dioxide  and  traces  of  ammonia,  is,  therefore, 
most  likely  to  dissolve  lead.  The  lead  eave-troughs,  etc.,  which  were 
once  extensively  used,  should,  therefore,  be  rejected,  in  case  the  rain- 
water is  used  for  drinking.  Well-water  usually  contains  acid  calcium 
carbonate  and  gypsum;  as  a  result,  the  lead  pipes  soon  become  coated 
.with  an  insoluble  layer  of  lead  sulphate  and  basic  carbonate  (as  well 
as  calcium  carbonate),  so  that  after  a  while  the  lead  can  no  longer  be 
absorbed  by  the  water. 

Lead  is  used  for  many  purposes,  not  only  in  the  elemental  con- 
dition, but  also  in  the  form  of  alloys  (see  §  199).  • 

Oxides  of  Lead. 

204.  The  following  oxides  of  lead  are  known:  Pb2O,  PbO, 
Pb2O3,  Pb304,  Pb02. 

Lead  oxide,  PbO,  is  the  only  one  of  these  oxides  with  basic  prop- 
erties. It  is  formed  by  direct  synthesis  from  its  elements  (§  203). 
It  is  fusible,  and  congeals  again  to  a  reddish-yellow  mass  called 
litharge.  By  carefully  heating  lead,  lead  hydroxide  or  lead  nitrate, 
the  oxide  is  obtained  as  an  amorphous  brown  powder  (massicot). 
It  is  somewhat  soluble  in  water,  forming  the  hydroxide.  It  dis- 
solves in  caustic  potash,  and  crystallizes  out  in  rhombic  prisms  on 
cooling. 


284  INORGANIC  CHEMISTRY.  [§§204- 

Lead  hydroxide,  Pb(OH)2,  is  formed  by  precipitating  a  lead 
solution  with  an  alkali.  It  is  amphoteric,  since  it  is  soluble 
in  caustic  alkalies;  ammonia,  however,  does  not  dissolve  it. 
On  being  warmed  to  145°  it  gives  up  water  and  turns  to 
oxide.  It  is  somewhat  soluble  in  water,  imparting  to  the  latter 
an  alkaline  reaction,  and  absorbs  carbon  dioxide  from  the 
air. 

Minium,  or  red  lead,  Pb3O4,  is  prepared  by  heating  lead  oxide 
or  white  lead  in  the  air  for  quite  a  while  at  300-400°.  Because 
of  its  pleasing  red  color  it  is  used  as  a  pigment  in  painting. 
Gentle  heating  makes  the  color  a  brighter  red  at  first;  stronger 
heating  turns  it  violet  and  finally  black;  on  cooling,  however, 
the  original  color  returns.  By  treating  it  with  dilute  nitric  acid, 
lead  nitrate  and  lead  peroxide  are  formed,  hence  minium  may  be 
regarded  as  2PbO-PbO2. 

Lead  peroxide,  PbO2,  is  obtained  in  the  way  just  stated;  more 
easily,  however,  by  passing  chlorine  into  an  alkaline  lead  solution 
or  adding  a  hypochlorite  to  a  lead  salt,  thus: 

2PbCl2  +  Ca(OCl)2  +  2H2O  =  2PbO2  +  CaCl2  +  4HC1. 

Milk  of  lime  is  then  added  to  neutralize  the  free  acid. 

Lead  peroxide  is  an  amorphous  dark-brown  powder.  It  has  the 
property,  common  to  most  peroxides,  of  giving  up  oxygen  easily. 
At  an  elevated  temperature  it  splits  up  into  lead  oxide  and  oxygen. 
On  warming  it  -with  sulphuric  acid,  lead  sulphate  and  oxygen  are 
formed;  on  warming  with  hydrochloric  acid,  lead  chloride  and 
chlorine  are  produced. 

Lead  peroxide,  like  the  oxides  C02  and  SnO2,  has  the  character 
of  an  acid  anhydride;  it  is  soluble  in  hot  concentrated  potassium 
hydroxide  and  this  solution,  on  cooling,  deposits  crystals  of  the  com- 
position K2PbOa  +  3H2O  (which  are  thus  entirely  analogous  in  com- 
position to  potassium  stannate) .  This  plumbate  is  easily  decomposed 
by  water  into  potassium  hydroxide  and  lead  peroxide.  If  we  regard 
lead  peroxide  as  an  acid  anhydride,  minium  can  be  considered  as 
the  lead  salt  of  the  normal  plumbic  acid,  Pb(OH)4,  i.e.  Pb2-Pb04. 
This  idea  is  confirmed  by  the  following  method  of  formation:  If  a 
solution  of  lead  oxide  in  potassium  hydroxide  is  added  to  a  solution 
of  the  plumbate  K2PbOs,  a  yellow  substance  is  precipitated  having 


205.]  HALOGEN  COMPOUNDS.  285 


the  composition  PbsC^-H^O,  which  gives  off  water  readily  and 
forms  minium. 

If  a  mixture  of  litharge  and  calcium  carbonate  is  heated  in  a  current 
of  air  at  700°,  carbon  dioxide  is  given  off,  oxygen  absorbed  and  cal- 
cium plumbate,  CaPbO3,  formed.  If  this  plumbate  is  treated 
with  carbon  dioxide  at  about  the  same  temperature,  calcium  carbonate 
and  lead  oxide  are  again  formed,  while  oxygen  escapes.  This  process 
(discovered  by  KASSNER)  serves  for  the  commercial  manufacture  of 
oxygen.  The  latter  is  brought  on  the  market  compressed  in  iron  bottles 
(c/.  also  §  262). 

The  oxide  Pb2O3  is  obtained  by  adding  sodium  hypochlorite  to 
a  solution  of  lead  oxide  in  potassium  hydroxide.  It  can  be  regarded 
as  the  lead  salt  of  a  lead  acid,  H2PbO3,  i.e.  as  Pb-PbOs,  for,  on 
treatment  with  dilute  nitric  acid  PbO2  (the  anhydride  of 
and  lead  nitrate  are  formed. 

Halogen  Compounds. 

205.  The  halogen  compounds  of  lead  having  the  formula 
are  difficultly  soluble  in  cold  water;  lead  fluoride  is  almost  insoluble 
and  the  solubility  of  the  three  others  decreases  with  increasing 
atomic  weight  of  the  halogen. 

Lead  chloride,  PbCl2,  is  obtained  as  a  white  precipitate  when 
dilute  hydrochloric  acid  is  added  to  the  solution  of  a  lead  salt.  At 
12.5°  it  dissolves  in  135,  at  100°  in  less  than  30,  parts  of  water 
and  crystallizes  from  the  hot  solution  in  the  form  of  white  silky 
needles  or  lamellse.  If  an  aqueous  solution  of  lead  chloride  is  treated 
with  dilute  hydrochloric  acid,  lead  chloride  is  precipitated,  for  by 
the  addition  of  Cl-ions  the  solubility  product  of  lead  chloride  is 
exceeded  ;  nevertheless,  lead  chloride  is  easily  soluble  in  concentrated 
hydrochloric  acid.  This  must  be  due  to  the  formation  of  a  com- 
pound of  the  chlorides  of  lead  and  hydrogen,  an  analogue  of  which 
has  been  found  in  PbI2-HI  +  10H2O,  which  has  been  isolated.  A 
characteristic  compound  of  lead  is  the  iodide  PbI2,  which  is  pre- 
cipitated from  lead  solutions  by  potassium  iodide.  It  is  scarcely 
soluble  in  cold,  but  moderately  soluble  in  hot  water.  It  crystallizes 
out  of  a  solution  in  dilute  acetic  acid  in  beautiful  crystal  flakes 
with  a  golden  lustre. 


2S6  INORGANIC  CHEMISTRY.  [§§206- 

Lead  tetrachloride,  PbCl4,  is  formed  when  a  solution  of  lead  dichloride 
in  strong  hydrochloric  acid,  cooled  by  ice2  is  saturated  with  chlorine. 
From  this  liquid  ammonium  chloride  precipitates  a  lemon-yellow  crys- 
talline substance,  2NH4C1  •  PbCl4,  having  a  composition  analogous  to 
pink-salt  (§  201). 

Analogous  double  salts  of  PbCl4  are  also  formed  with  the  alkali 
metals,  such  as  potassium  and  rubidium  (Rb2PbCl6).  If  one  of  these 
double  salts  is  treated  with  ice-cold  concentrated  sulphuric  acid,  lead 
tetrachloride  gradually  separates  out  as  a  heavy  yellow  pil  (sp.  g.  3.18), 
which  is  stable  at  a  low  temperature  and  crystallizes  at  —15°.  At  as 
high  a  temperature  as  room  temperature,  and  more  rapidly  on  warming, 
it  breaks  up  into  lead  dichloride  and  chlorine. 

Other  Lead  Salts. 

206.  Lead  nitrate,  Pb(N03)2,  is  obtained  by  dissolving  lead  in 
dilute  nitric  acid.  It  is  colorless,  crystallizes  isometric  and  is 
soluble  in  8  parts  of  water.  Heating  decomposes  it  (§  122). 
Several  basic  lead  nitrates  are  known 

Lead  sulphate,  PbSCU,  is  practically  insoluble  in  water  and  can 
therefore  be  obtained  by  precipitating  a  lead  solution  with  dilute 
sulphuric  acid  or  a  soluble  sulphate.  It  occurs  as  a  mineral  in 
crystallized  form  under  the  name  of  anglesite;  it  is  isomorphous 
with  the  sulphate  of  barium,  barite.  Lead  sulphate  is  soluble  in 
concentrated  sulphuric  acid;  hence  the  crude  acid  which  is  con- 
centrated in  lead  pans  (§  186,  3)  contains  lead  sulphate;  this  is 
precipitated  on  diluting  the  acid  with  water.  It  is  dissolved  by 
concentrated  alkalies.  Ignition  on  charcoal  reduces  it  to  sulphide. 
Lead  disulphate,  " plumbic  sulphate,"  Pb(SO4)2,  separates  from 
the  acid  around  the  anode  when  sulphuric  acid  of  1.7-1.8  specific 
gravity  is  elect roly zed  between  lead  electrodes.  It  has  not  been 
obtained  quite  free  from  lead  sulphate.  It  is  a  white  granular  sub- 
stance of  strong  oxidizing  properties.  Water  decomposes  it  readily 
into  sulphuric  acid  and  lead  peroxide.  It  is  isomeric  with  lead 
persulphate,  PbS2Og,  a  salt  of  the  dibasic  persulphuric  acid. 

Lead  carbonate,  PbC03,  is  deposited  when  a  solution  of  the 
nitrate  is  treated  with  ammonium  carbonate.  White  lead,  a  basic 
carbonate,  is  used  extensively  as  a  pigment.  However,  it  soon 
turns  black,  if  any  hydrogen  sulphide  (from  drainage  pipes,  etc.) 
comes  in  contact  with  it;  moreover,  it  is  injurious  to  the  health, 
because  it  comes  off  of  the  painted  walls  in  the  form  of  dust  and 


207.] 


SUMMARY  OF  THE  CARBON  GROUP. 


287 


gets  into  the  lungs.  White  lead  is  particularly  valuable  for  its 
covering-power,  i.e.  the  painted  surface  appears  perfectly  white 
when  covered  with  only  a  very  thin  layer  of  the  pigment;  it  is 
much  greater  than  that  of  other  white  pigments,  such  as  white  zinc 
and  barite,  which  are  frequently  substituted  for  white  lead  because 
they  are  harmless. 

The  manufacture  of  white  lead  is  still  carried  on  exten- 
sively after  the  Dutch  method.  This  consists  in  placing  rolls  or 
"buckles"  of  lead-plate  into  jars  containing  a  little  acetic  acid.  The 
vessels  are  loosely  covered  with  a  leaden  lid  and  buried  in  a  heap  of 
horse-manure.  The  heat  generated  by  the  decaying  manure  causes 
a  part  of  the  acetic  acid  to  evaporate  and  converts  the  lead  into  basic 
lead  acetate.  The  latter  is  then  transformed  to  white  lead  by  the  car- 
bon dioxide  given  off  from  the  decaying  heap.  After  about  five  or 
six  weeks  the  plates  are  almost  entirely  changed  to  white  lead.  This 
is  then  ground  moist,  washed  out  (to  remove  any  acetate)  and  dried 
whereupon  it  is  sent  to  the  market. 

Lead  sulphide,  PbS,  is  black  and  comes  down  amorphous  when 
hydrogen  sulphide  is  passed  into  a  lead  solution.  A  liquid  con- 
taining only  traces  of  lead  is  colored  brown  by  sulphuretted  hy- 
drogen; this  is  a  very  delicate  means  of  testing  for  lead.  Strong 
nitric  acid  oxidizes  it  readily  to  lead  sulphate. 


SUMMARY  OF  THE   CARBON   GROUP. 

207.  The  elements  carbon,  silicon,  germanium,  tin  and  lead 
form  a  natural  group,  as  may  be  seen  from  a  comparison  of  their 
physical  and  chemical  properties.  In  the  following  table  the 
most  important  physical  constants  are  summarized ;  here, 
as  in  other  natural  groups,  the  gradual  change  of  these  constants 
with  the  rise  of  the  atomic  weight  is  very  evident: 


C. 

Si. 

Ge. 

Sn. 

Pb. 

Atomic  weight. 

12.00 

28.3 

72.5 

119.0 

207.10 

Specific  gravity. 

f    2.251 
1    3.6    ) 

2.49 

5.5 

7.29 

11.39 

Melting-point.  . 
Boiling-point.  .  . 

above  3000° 

1420° 

circa  900° 

233° 
circa  1500° 

327° 
1525° 

288  INORGANIC  CHEMISTRY.  [§§207- 

With  respect  to  the  chemical  properties  we  note  in  the 
first  place  that  all  these  elements  have  the  same  compound-types, 
MX2  and  MX4;  in  other  words,  that  they  are  bi-  or  quadri-valent; 
this  is  even  true  of  lead  (PbO2,  PbCl4,  etc.),  which  does  not  fit 
into  the  table  of  physical  properties  with  its  boiling-  and  melting- 
points.  Moreover,  there  is  to  be  noted  in  general  a  transition 
from  metalloid  to  metallic  character,  as  is  plainly  shown  by  the 
following  facts: 

1.  Only  carbon  and  silicon  are  known  to  form  hydrogen  com- 
pounds (of  an  indifferent  nature). 

2.  Of  the  oxygen  compounds  of  the  MO  type,  that  of  carbon  is 
indifferent  and  the  others  (no  such  compound  of  silicon  is  known) 
grow  more  ba^ic  in  character  as  the  atomic  weight  increases,  lead 
hydroxide  having  rather  strongly  alkaline  properties. 

3.  The  oxygen  compounds,  RO2,  however,  are  decidedly  acidic 
in  character  in  the  cases  of  carbon  and  silicon  and  also  in  the  case 
of  germanium,  while  in  that  of  lead  the  salts  of  the  acid  H2PbO3 
are  immediately  decomposed  by  water,  so  that  here  the  acid  proper- 
ties appear  much  weakened. 

4.  As  to  the  halogen  compounds,  those  of  carbon  (CX4)  are 
unaffected  by  cold  water — perhaps  because  of  their  insolubility 
in  it;    the  other  halogen  compounds,  MX4,  are  decomposed  by 
water. 

Lead,  in  some  of  its  physical  and  chemical  properties,  does  not  dis- 
play the  gradation  which  is  ordinarily  met  with  in  the  elements  of  a, 
group.  This  phenomenon  is  quite  often  observed  in  elements  of  very 
high  atomic  weight.  In  the  nitrogen  group  we  saw  it  in  the  case  of 
bismuth. 

METHODS  OF  DETERMINING  ATOMIC  WEIGHTS. 

208.  So  far  only  one  method  of  determining  the  atomic  weight 
has  been  mentioned  (§  34).  This  consists  in  investigating  as  large 
as  possible  a  number  of  gaseous  compounds  of  the  element  as  to 
their  vapor  density  and  empirical  composition  and  then  calculating 
how  many  grams  of  the  element  are  contained  in  a  mole  of 
the  various  compounds.  The  smallest  figure  thus  found  is  taken 
as  the  atomic  weight.  Although  this  method  is  quite  general, 


208.J       METHODS  OF  DETERMINING   ATOMIC   WEIGHTS.       289 


it  has  the  drawback  of  affording  only  a  certain  degree  of  proba- 
bility, a  probability  which  becomes  greater  as  the  number  of  in- 
vestigated compounds  increases  and  which  lessens  the  chance  of 
finding  a  compound  that  contains  per  mole  only  a  simple  fraction 
of  the  previously  accepted  atomic  weight. 

There  are,  however,  other  methods.  None  of  them  is  so 
generally  applicable  as  this,  but  they  are  of  a  more  absolute  char- 
acter and  have  been  of  great  value  in  the  many  cases  in  which 
they  could  be  used.  They  furnish  a  very  valuable  check  on  the 
determinations  made  by  the  general  method.  These  methods  are 
based  on  the  following  laws : 

1.  The  law  of  DULONG  and  PETIT.  The  product  of  the  atomic 
weight  of  a  s  oli  d  element  and  its  specific  heat  is  about  6.4.  This 
is  evident  from  the  table  on  the  opposite  page. 

Most  of  the  values  of  the  product  lie  as  the  table  shows,  very 
close  to  6.4;  the  maximum  is  6.9,  the  minimum  5.0.  Calling  this 
product  the  atomic  heat,  we  can  express  the  law  of  DULONG  and 
PETIT  in  the  following  simple  way :  The  atomic  heat  of  the  solid  ele- 
ments is  approximately  constant  and  is  about  6.4. 

A  few  deviations  have  been  pointed  out  in  the  last  column;  in 
such  cases  it  has  frequently  been  found,  however,  that  at  an 
increased  temperature  the  atomic  heat  approaches  the  value  6.4. 
This  is  probably  due  to  the  fact  that  at  the  temperature  (room- 
temperature)  at  which  the  measurements  of  the  specific  heat  of 
the  elements  have  been  mainly  carried  out,  the  elements  are  not 
all  in  the  proper  physical  condition  for  comparison.  It  is  notably 
the  elements  with  atomic  weights  under  35  that  show  the  greatest 
deviations. 

It  is  an  interesting  fact  that  there  is  a  certain  regularity  to  be  found 
in  these  irregularities.  The  latter  become  more  marked  as  the  valence 
increases. 


Element  .....  Li 

Valence  ......   1 

Atomic  heat..  6.  6 


Be       B 
23 

3.7    2.8 


C 

4 

1.9 


Na 

1 

6.7 


Mg 
2 

6.1 


Al 
3 

5.8 


Si 

4 

4.6 


P        S 
3(5)  2(4,6) 
5.9      5.7 


It  is  easy  to  see  how  the  law  of  DULONG  amd  PETIT  can  be 
made  use  of  for  the  determination  of  atomic  weights.     Inverting, 

we  have  -   -—  =At.  wt. 

b.   H. 


290 


INORGANIC  CHEMISTRY. 


[§  209- 


Element. 


Sp.  H. 


At.  Wt. 


Product. 


Remarks. 


Hydrogen. 
Lithium.  . , 
Beryllium. 


Boron. 


^    ,       f  amorphous. 
Carbon    diamond. 


Sodium.  . . . 
Magnesium. 
Aluminium. 
Silicon.  . 


Phosphorus. 


Sulphur. 
Potassium. . . 

Calcium 

Scandium.  . . 
Chromium. . . 
Manganese. . . 

Iron 

Cobalt 

Nickel 

Copper 

Zinc 

Gallium 

Germanium. . 

Arsenic 

Selenium.  . . . 

Bromine 

Zirconium.  . . 
Molybdenum. 
Ruthenium.  . 
Rhodium. . . . 
Palladium.  . . 

Silver 

Cadmium. . . . 

Indium 

Tin. 

Antimony. .  . 
Tellurium.  . . , 

Iodine 

Lanthanum. . 

Cerium 

Tungsten. .    . 

Osmium 

Iridium 

Platinum.  . . . 

Gold 

Mercury 

Thallium.  . . . 

Lead 

Bismuth 

Thorium.  . . . 
Uranium.  .  . . 


6 

0.941 

0.408 

0.254 

0.174 
0  143 

0.293 
0.250 
0.214 
0.165 

0.189 

0.178 
0.166 
0.170 
0.153 
0.121 
0.122 
0.114 
0.107 
0.108 
0.095 
0.094 
0.079 
0.077 
0.082 
0.080 
0.084 
0.066 
0.072 
0.061 
0.058 
0.059 
0.057 
0.054 
0.057 
0.054 
0.051 
0.047 
0.054 
0.045 
0.045 
0.033 
0.031 
0.032 
0.032 
0.032 
0.032 
0.033 
0.031 
0.030 
0.027 
0.027 


1.008 

6.94 

9.1 

11.0 
I  12. 00 

23.00 
24.32 
27.1 

28.3 

31.04 

32.07 
39.10 
40.09 
44.1 
52.0 
54.93 
55.85 
58.97 
58.68 
63.57 
65.37 
69.9 
72.5 
74.96 
79.2 
79.92 
90.6 
96.0 
101.7 
102.9 
106.7 
107.88 
112.40 
114.8 
119.0 
120.2 
127.5 
126.92 
139.0 
140.25 
184.0 
190.9 
193.1 
195.2 
197.2 
200.0 
204.0 
207.10 
208.0 
232.4 
238.5 


6 

6.6 

3.7 

2.8 

2.1 
1.7 

6.7 
6.1 
5.8 
4.6 

5.9 

5.7 
6.5 
6.8 
6.7 
6.3 
6.7 
6.4 
6.3 
6.4 
6.0 
6.1 
5.5 
5.6 
6.9 
6.3 
6.7 
6.0 
6.9 
6.3 
6.0 
6.0 
6.1 
6.0 
6.5 
6.5 
6.1 
6.0 
6.8 
6.2 
6.3 
6.1 
5.9 
6.1 
6.2 
6.3 
6.4 
6.7 
6.4 
6.2 
6.2 
6.5 


Liquid. 

Sp.  H.  at  257° =0.58;  prod. 

=  5.2. 
Amorphous.    Sp.  H.  at  400° 

=  0.58;  prod.  =6.4. 

Above  900°  Sp.  H.=  0.459; 
prod.  =5.5. 


Crystallized.  Sp.  H.  above 
200°  =  0.204;  prod.  =5.8. 

Yellow.  Sp.  H.  of  redP- 
0.1698;  prod.  =5.24. 

Rhombic. 


Crystallized 

Do. 

Solid. 


Solid 


209.]       METHODS  OF  DETERMINING  ATOMIC  WEIGHTS.        291 

Of  course  the  result  thus  obtained  is  only  approximately 
correct,  for  the  product  6.4  is  not  strictly  constant.  The  method 
is,  however,  reliable  enough  to  determine  what  multiple  of  the 
equivalent  weight  (§  22),  the  exact  value  of  which  can  be  found 
by  analysis,  is  the  atomic  weight. 

209.  2.  Closely  connected  with  the  law  just  enunciated  is  that 
of  NEUMANN",  which  has  been  more  carefully  investigated  by  REG- 
NAULT  and  KOPP.  This  law  says  that  in  solid  compounds  each 
element  has  a  constant  atomic  heat,  which  varies  but  little  from  that 
of  the  free  element.  The  molecular  heat  is  therefore  equal  to  the 
sum  of  the  atomic  heats.  If  the  molecular  heat  of  solid  com- 
pounds is  divided  by  the  number  of  atoms,  the  quotient  must  be 
about  6.4.  In  reality  this  quotient  proved  to  be:  for  bromine 
compounds  RBr,  6.9;  for  RBr2,  6.5;  for  iodine  compounds  RI,  6.7; 
RI2,  6.5.  The  law  of  NEUMANN  likewise  holds  for  many  elements 
whose  specific  heat  in  the  solid  state  has  not  been  susceptible  of 
measurement,  thus  e.g.  for  chlorine  compounds:  for  RC1  com- 
pounds the  quotient  referred  to  was  6.4,  for  RC12,  6.2,  for  certain 
double  chlorides,  6.1-6.2.  For  other  elements,  like  oxygen,  the 
atomic  heat  found  from  the  molecular  heat  of  the  compounds  is 
constant,  but  it  is  about  4.0  instead  of  about  6.4.  The  same  is 
true  of  hydrogen,  whose  mean  atomic  heat  in  solid  compounds 
is  2.3.  These  figures  were  found  by  determining  the  molecular 
heat  of  various  oxygen  or  hydrogen  compounds  and  subtracting 
from  it  the  known  atomic  heats  of  the  other  elements.  If  the 
atomic  heat  thus  obtained  is  divided  by  the  atomic  weight,  we 
have  the  specific  heat  of  the  element  in  its  compounds. 

The  way  in  which  the  law  of  NEUMANN  can  be  applied  to 
atomic  weight  determinations  is  illustrated  by  the  following 
example: 

The  problem  is  to  determine  the  atomic  weight  of  calcium  with  the 
help  of  the  specific  heat  of  sulphate  of  lime,  CaS04,  which  amounts  to 
0.2  according  to  REGNAULT. 

Analysis  has  shown  that  anhydrous  calcium  sulphate  contains  29.4% 
calcium,  23.5%  sulphur,  and  47.1%  oxygen.  Since  sulphates  contain  the 
SO4  group  in  combination  with  a  metal,  it  follows  from  the  above 
.analysis  that  there  must  be  associated  with  this  group  40  parts  by 
weight  of  calcium.  The  next  question  is,  whether  40  is  really  the 


292  INORGANIC  CHEMISTRY.  [§§  209- 

atomic  weight  of  calcium  or  a  multiple  or  submultiple  of  the  atomic 
weight. 

Now  the  molecular  weight  of  calcium  sulphate  must  be  40  +  32  +4 X 16 
=  136,  no  matter  whether  40  is  the  relative  weight  of  one  or  of  more 
than  one  calcium  atom.  The  molecular  heat  is  therefore  136  X  0.2  =27.2. 
The  atomic  heat  of  sulphur  in  compounds  is  about  5.4  and  that  of  oxygen 
about  4.0;  consequently  the  molecular  heat  of  the  SO4  group  in  its  solid 
compounds  is  5.4+4X4.0  =  21.4.  For  the  atomic  heat  of  calcium  we 
have  the  remainder,  27.2-21.4  =  5.8. 

It  therefore  follows  that  the  formula  of  calcium  sulphate  must  be 
CaS04,  which  means  that  40  is  the  atomic  weight  of  calcium;  for,  if  the 
atomic  weight  were  a  multiple  or  submultiple  of  this  number,  we  should 
have  found  for  the  atomic  heat  of  the  metal  a  number  much  farther 
from  the  average  atomic  heat  of  the  elements,  6.4,  than  5.8. 

The  value  of  the  atomic  weight  calculated  from  NEUMANN'S 
law  therefore  serves  merely  to  decide  what  multiple  of  the  equiva- 
lent weight  must  be  taken ;  for  this  purpose  the  number  so  obtained 
is  sufficiently  accurate. 

210.  3.  The  law  of  Mitscherlich.  The  crystal  form  of  com- 
pounds having  analogous  chemical  composition  is  the  same:  or,  in 
other  words,  compounds  of  analogous  chemical  composition  are 
is  omor  phous .  The  compounds  KC1,  KI,  KBr,  e.g.  are 
analogous  in  composition;  they  all  crystallize  in  cubes.  H2KP04, 
H2KAsC>4,  H2(NH4)P04  also  have  an  analogous  composition  and 
all  crystallize  in  the  tetragonal  system.  The  analogous  compounds 
KC104  and  KMn(>4  both  crystallize  rhombic. 

If  two  compounds  have  been  proved  to  be  isomorphous,  it  is 
very  probable  that  their  composition  is  analogous,  whereupon 
the  atomic  weight  is  readily  found.  Let  us,  for  example,  take 
the  case  of  manganese,  supposing  its  atomic  weight  to  be  unknown; 
now  potassium  permanganate  is  isomorphous  with  potassium  per- 
chlorate,  which  latter  is  known  to  have  the  formula  KC104. 
Analysis  has  shown  the  formula  of  potassium  permanganate  to  be 
KMnx04,  x  being  unknown,  for  39  parts  (by  weight)  of  potassium 
(1  atom)  are  combined  with  64  parts  of  oxygen  (4  atoms)  and  55 
parts  of  manganese  (x  atoms).  From  its  isomorphism  with  KCICU 
it  follows  that  its  formula  must  be  KMn(J4  (i.e.  x  =  l),  hence  55 
is  the  atomic  weight  of  manganese. 


211.]          DETERMINATION  OF  EQUIVALENT  WEIGHTS.  293 

In  determining  the  atomic  weight  of  zinc  we  could  use  the  iso- 
morphism of  the  crystallized  sulphates  of  magnesium  and  zinc.  The 
formula  of  the  former  is  MgSO44-7H2O.  On  the  basis'  of  the 
analysis  of  zinc  sulphate  and  the  isomorphism  mentioned  we  have 
the  formula  ZnS04+7H2O,  from  which  the  atomic  weight  is  ob- 
tained in  the  same  way  as  above. 

The  law  of  isomorphism  was  discovered  as  early  as  1819.  Since  at  that 
time  the  law  of  AVOGADRO  received  little  attention  and  the  determination 
of  the  specific  heat  was  in  many  cases  impossible,  the  phenomena  of  iso- 
morphism were  the  most  important  means  of  getting  information  regarding 
the  value  of  the  atomic  weight.  Subsequently  its  importance  for  this  pur- 
pose lessened,  mainly  because  simpler  means  were  found,  but  also  because 
it  proved  to  be  very  difficult  in  many  cases  to  decide  whether  two  sub- 
stances are  isomorphous.  Moreover  it  was  found  that  certain  substances 
of  entirely  different  composition  are  isomorphous. 

A  very  delicate  test  for  isomorphism  is  the  fact  that  a  supersaturated 
solution  can  be  made  to  crystallize,  not  only  by  an  extremely  small  amount 
of  the  dissolved  substance  itself  ("  sowing,"  or  "  inoculation "),  but  by  bodies 
that  are  isomorphous  with  it. 


Experimental  Determination  of  Equivalent  Weights. 

2ii.  In  the  methods  above  described  the  question  is  one  of  determining 
"which  multiple  of  the  equivalent  weight  is  the  atomic  weight.  In  order 
to  establish  the  atomic  weights  with  accuracy  the  equivalent  weights  must 
be  determined  with  the  greatest  possible  precision.  The  solution  of  this 
problem,  which  is  one  of  fundamental  importance,  since  all  the  numerical 
relationships  of  chemical  reactions  are  based  on  the  atomic  weights,  has 
been  the  object  of  numerous  investigations  in  the  preceding  century  and 
to-day  it  is  still  only  partially  accomplished. 

The  first  atomic  weight  table  dates  from  DALTON  in  1805.  The  figures 
given  in  it  were  scarcely  more  than  rough  approximations.  BERZELIUS 
(1779-1848)  in  the  first  and  second  decades  of  the  century  determined  a 
long  series  of  equivalent  numbers,  after  having  been  first  obliged  in  most 
cases  to  work  out  reliable  analytical  methods.  The  atomic  weights  at 
which  he  arrived  were  in  general  use  for  many  years  and  really  differ  from 
the  more  accurate  ones  now  employed  by  hardly  more  than  a  fraction  of  a 
per  cent.  Exceedingly  accurate  "atomic  weight  determinations"  were 
undertaken  by  STAS  (1813-1891).  The  ten  atomic  weights  determined  by 
him,  viz.  those  of  Ag,  Cl,  Br,  I,  K,  Na,  Li,  S,  Pb,  and  N,  are  in  most  cases 
accurate  to  within  a  few  units  in  the  second  decimal  place.  The  researches 
called  for  most  exhausting  and  persistent  labors  during  a  long  period  of 
years. 


294  INORGANIC   CHEMISTRY.  [§§211- 

In  the  last  decade  atomic  weight  determinations  have  been  carried 
out  on  a  scale  of  much  greater  refinement  by  MORLEY,  RICHARDS,  GUYB 
and  others  and  the  accuracy  of  the  values  has  been  extended  another  decimal 
place,  so  that  now  not  a  few  of  the  atomic  weights  are  established  with 
certainty  to  within  a  few  units  in  the  third  decimal  place. 

In  determining  atomic  weights  either  purely  chemical  or  physico- 
chemical  methods  may  be  employed.  Both  have  been  greatly  perfected 
in  these  latter  investigations  and  they  will  now  be  described  in  a  few  para- 
graphs. 

As  for  the  purely  chemical  methods,  there  are  four  conditions  which 
are  essential  to  an  accurate  determination  of  an  atomic  weight:  (a)  A 
suitable  substance  must  be  found  which  can  be  prepared  perfectly  pure. 
(6)  This  compound  must  contain  in  addition  to  the  element  under  study 
only  elements  of  accurately  known  atomic  weight,  (c)  The  valence  of  the 
elements  in  this  compound  must  be  well  denned.  It  is  not  permissible, 
for  example,  that  the  substance  be  a  mixture  of  two  stages  of  oxidation. 
(d)  The  compound  selected  must  be  adapted  to  an  exact  analysis,  or  else 
its  exact  synthesis  from  the  weighed  elements  must  be  possible. 

Notwithstanding  the  simplicity  and  legitimacy  of  these  demands  it  is 
often  difficult  to  satisfy  them.  The  preparation  of  a  compound  in  the  pure 
state  is  among  the  most  difficult  of  operations,  if  by  purity  we  mean  the 
reduction  of  the  impurities  to  a  10 ~4  part  of  the  whole.  It  was  formerly 
believed  that  this  could  be  readity  accomplished  by  recrystallization,  but 
now  we  know  that  every  substance  that  separates  out  in  a  solid  phase  has 
a  tendency  to  retain  upon  its  surface  or  in  its  interior  a  part  of  the  other 
substance  contained  in  the  phase  out  of  which  the  solid  separated.  All 
precipitates  or  crystals  from  aqueous  solutions  contain  water  that  is  not 
in  chemical  combination.  Even  the  splendid  glistening  silver  crystals  that 
are  obtained  in  the  electrolysis  of  a  silver  nitrate  solution  and  are  apparently 
perfectly  dry  and  pure  contain  not  only  water  but  silver  nitrate  as  well. 
Silver  chloride,  precipitated  from  a  solution  of  sodium  chloride  by  silver 
nitrate,  may  have  included  traces  of  NaCl,  AgNO3  or  NaNO3,  even  after  a 
thorough  washing.  Potassium  chlorate,  though  much  less  soluble  than 
potassium  chloride,  contains  nevertheless  0.027%  of  the  latter  after  repeated 
recrystallizations.  One  of  the  most  troublesome  sources  of  error  in  all 
quantitative  researches  is  the  unsuspected  presence  of  hygroscoDically 
held  water,  since  it  is  not  at  all  easy  to  detect  by  chemical  tests  and  causes 
no  essential  change  in  the  external  appearance  of  the  substance  containing 
it. 

The  analysis  of  a  substance  resolves  itself  in  most  cases  into  a  separation 
cf  its  components  in  the  form  of  other  compounds  and  weighings  of  the 
latter.  For  example,  in  order  to  determine  the  silver  content  of  silver 
nitrate  the  metal  is  thrown  down  as  silver  chloride  and  the  latter  is  weighed, 
whereupon  the  quantity  of  silver  can  be  calculated  from  the  known  silver 
content  of  the  chloride.  The  analyst  generally  finds  it  also  necessary  to 
convert  one  compound  into  another  quantitatively.  The  modern  investiga- 


212.]          DETERMINATION  OF  EQUIVALENT  WEIGHTS.  295 

tions  of  atomic  weights  have  also  taught  us  that  this  is  often  a  very  difficult 
problem.  Among  other  sources  of  error  in  this  connection  are  the  solu- 
bility of  the  so-called  "  insoluble"  substances  and  the  solubility  of  glass. 
It  has  long  been  known  that  substances  like  silver  chloride,  barium  sul- 
phate, etc.,  are  not  strictly  insoluble;  but  their  solubility  has  first  received 
proper  attention  in  connection  with  the  recent  atomic  weight  determinations. 
In  working  with  glass  vessels  it  is  impossible  to  avoid  silicic  acid  as  an 
impurity.  Recognition  of  this  fact  has  led  to  the  use  of  vessels  of  quartz 
or,  better  still,  of  platinum,  which  has  proved  to  be  an  important  refinement 
of  method. 

212.  Physico-chemical  methods  have  found  application  in  the  determination 
of  the  volume  weight  of  gases.  One  of  the  most  fruitful  of  modern  physical 
concepts  is  that  of  the  ideal  gas,  whose  expansion  at  constant  pressure  or 
pressure  increase  at  constant  volume  both  have  a  coefficient  for  a  tempera- 
ture change  of  one  degree  of  exactly  1/273.08.  Moreover,  the  ideal  gas 
is  in  strict  accord  with  BOYLE'S  law.  A  gram  molecule  of  such  a  gas  at 
6°  and  760  mm.  Hg  pressure  would  occupy  a  volume  of  22.412  1.  However, 
the  actual  gases  are  more  compressible  and  expansible  than  the  ideal  gas; 
hydrogen  and  helium  are  the  only  ones  that  are  less  compressible.  For  this 
reason  22.412  1.  of  an  actual  gas  at  0°  and  760  mm.  Hg  contains  a  little 
more  than  one  gram  molecule.  If  we  let  1  +  A  represent  the  number  of 
gram  molecules  of  an  actual  gas  which  are  contained  in  22.412  1.,  the  molecular 
weight  of  the  gas  becomes 

22.412  G 


where  G  is  the  weight  of  a  liter  of  the  gas  under  normal  conditions.  The 
establishment  of  the  atomic  weight  of  a  gas  thus  resolves  itself  into  the 
accurate  determination  of  the  magnitudes  G  and  A.  The  methods  for 
ascertaining  the  exact  weight  of  a  given  volume  of  a  gas  have  undergone 
important  improvements  in  recent  years.  The  agreement  of  the  values 
found  by  the  different  investigators  is  within  ±0.0001.  While  formerly 
the  gases  were  weighed  in  huge  globes,  some  containing  as  much  as  21  1., 
later  investigators  have  been  able  to  reduce  this  volume  to  between  one  liter 
and  half  a  liter,  or  even  less.  Nevertheless  the  concordance  between  the 
various  series  of  determinations  was  improved,  because  the  corrections  for 
the  small  globes  were  much  less.  An  additional  correction  was  applied 
for  the  contraction  of  a  globe  on  evacuation,  due  to  the  external  pressure 
of  the  atmosphere  reducing  the  volume  slightly;  the  buoyancy  effect  of  the 
air  is  somewhat  less  for  an  evacuated  globe  than  for  one  filled  with  gas. 

In  order  to  remove  completely  the  layer  of  air  that  has  been  condensed 
on  the  inner  surface  of  the  globe,  it  is  necessary  to  evacuate  the  latter 
repeatedly  to  as  low  a  vacuum  as  possible  and  to  fill  it  with  the  gas  whose 
density  is  to  be  determined,  great  care  being  taken  meanwhile  to  exclude 
the  air. 

Furthermore,  the  purification  of  the  gases  to  be  weighed  is  much  better 


296  INORGANIC  CHEMISTRY.  [§§  212- 

accomplished  by  first  liquefying  them  and  then  removing  the  impurities 
by  fractional  distillation  at  a  lo\?  temperature.  ' 

The  determination  of  the  quantity  A  can  be  accomplished  in  four  different 
ways,  which  are  found  described  in  the  larger  physics  manuals.  It  should 
be.  noted,  however,  that  they  have  not  yet  attained  the  exactness  that 
characterizes  the  methods  of  determining  G.  In  the  cases  of  the  less  easily 
condensed  gases,  like  H,  N,  O,  and  Cl,  however,  very  accurate  determinations 
have  already  been  made. 

From  an  experimental  standpoint  these  physico-chemical  methods  have 
a  decided  superiority  over  the  purely  chemical  methods  in  that  ph^Gica' 
measurements  only  are  carried  out  after  the  gas  has  been  obtaiivxi  pure, 
All  the  uncertainties  that  are  involved  in  chemical  transformations  are  thus 
avoided;  and  upon  such  transformations  every  purely  chemical  determi- 
nation of  an  atomic  weight  is  based. 


THE  PERIODIC  SYSTEM  OF  THE  ELEMENTS. 

213.  In  studying  the  elements  which  we  have  considered  so  far, 
we  have  found  that  they  can  be  arranged  into  groups  of  elements 
according  to  their  valence,  the  elements  of  each  group  showing 
great  similarity  in  the  types  of  their  compounds.  The  physical 
and  chemical  properties  of  the  elements  of  such  a  group  are 
found  to  change  progressively  as  the  atomic  weight  increases. 
The  question  now  arises  whether  all  elements  can  be  thus  arranged 
into  groups;  the  reply  is  affirmative. 

In  the  course  of  the  last  century  there  was  no  lack  of  attempts  to  arrange 
the  elements  into  groups  of  similar  elements.  DOEBEREINER  called  atten- 
tion to  a  simple  relation  between  the  atomic  weights  of  kindred  elements 
as  early  as  1817,  and  in  1829  he  presented  the  doctrine  of  triads,  i.e.  he 
showed  that  there  are  different  groups  of  three  elements  each,  which  have 
a  great  similarity  among  themselves  and  a  constant  difference  in  the  atomic 
weights,  e.g.  Cl,  Br,  I;  Ca,  Sr,  Bat  etc.  In  the  year  1865  the  law  of  octaves 
was  proposed  by  NEWLANDS,  he  having  discovered  that,  if  the  elements 
are  arranged  according  to  increasing  atomic  weight,  after  an  interval  of 
seven  elements  an  element  follows  which  has  properties  analogous  to  those 
of  the  first,  i.e.  the  first,  eighth,  fifteenth,  etc.,  are  similar.  In  1869 
MENDELEEFF  and  LOTHAR  MEYER  almost  simultaneously  reached  con- 
clusions which  are  comprehended  by  the  term  "  periodic  system." 

If  we  arrange  the  elements  according  to  increasing  atomic 
weight,  thus: 

H  1 

Li  7        Be    9.1       B  11         C  12         N  14         O  16      F  19 
Na23      Mg24         Al  27.1      Si  28.4      P  31.0      S  32      Cl  35.4 


213.]  THE  PERIODIC   SYSTEM  OF  THE  ELEMENTS.        297 

we  see  that  there  is  a  gradual  variation  in  the  properties  of  elements 
in  a  horizontal  line;  after  fluorine,  however,  a  small  increase  in  the 
atomic  weight  involves  a  sudden  change  of  properties.  Moreover 
those  elements  which  are  in  the  same  vertical  column  show  great 
similarity,  as  we  saw  above  in  the  cases  of  carbon  and  silicon, 
nitrogen  and  phosphorus,  etc. 

This  regular  change  makes  itself  evident  in  the  valence  toward 
oxygen,  which  rises  from  one  (with  Li  and  Na)  to  two  (Be,  Mg), 
hree  (B,  Al),  four  (C,  Si),  five  (N,  P),  six  (S)  and  seven  (Cl  in  C12O7). 
The  valence  toward'  hydrogen v or  a 'halogen  increases,  however, 
from  one  (Li)  to  four  (C)  and  then  falls  again  to  one  (F).  A  similar 
regular  change  is  to  be  observed  with  reference  to  the  physical 
properties,  e.g.  specific  gravity  and  atomic  volume. 

Na          MR          Al  Si         P  (red)         S       Cl  (liq.) 

Sp.  gravity 0.97     1.75    2.67    2.49    2.14    2.06     1.33 

AT.  volume 24        14        10        11        14        16         27 

by  atomic  volume  we  understand  the  atomic  weight  divided  by  the 
weight  of  the  unit  volume  (based  on  water  of  4°  as  1);  it  is  therefore 
tLe  number  of  cubic  centimeters  occupied  by  a  gram-atom. 

Here  we  observe  an  increase  of  the  specific  gravity  up  to  alu- 
minium, then  again  a  decrease  to  chlorine,  while  the  atomic  vol- 
ume, on  the  other  hand,  decreases  from  the  beginning  of  the  series 
to  aluminium  and  then  increases.  This  steady  change  of  the 
same  physical  properties  is  also  observed  in  the  compounds 
of  the  above  elements.  For  the  oxides,  e.g.  we  have: 

Na2O       M-rO      Al2Oa       SiO2       P2O5        SO3      Cl2Or 

Sp.  gravity 2.8      3.7      4.0      2.6      2.7       1.9       ? 

At.  volume 22        22        25        45        55        82        ? 

Moreover,  if  we  write  down  a  series  of  elements  according  to 
increasing  atomic  weight,  beginning  with  another  univalent  metal, 
we  discover  irregularities  of  exactly  the  same  sort  as  the  above. 
The  following  series  may  serve  as  an  example  of  this: 

Ag  Cd  In  Sn  Sb  Te  I 

Atomic  wt 107.8     112.4      115       119.0    120.2     127.6     127.0 

Sp.  gravity 10.5        8.6      7.4          7.2        6.7        6.2        4.9 

Heie  also  we  find  the  same  gradual  rise  of  valence  from  silver, 
which  is  univalent,  to  septivalent  iodine,  the  progressive  transi- 


298  INORGANIC  CHEMISTRY.  [§   213. 

tion  from  metal  to  metalloid  and  a  continuous  decrease  in  specific 
gravity.  But  more;  if  we  put  this  last  row  under  the  first  two: 

H    1 

Li       7       Be      9.1    B      11       C      12       N     14       0.     16       F     19 
Na   23.1    Mg   24.4    Al    27.1    Si     28.4    P      31.0    S      32.1    Cl    35.5 
Ag  107.9    Cd  112.4    In  115       Sn  119.0    Sb  120.2    Te  127.6    I    127.0, 

it  is  apparent  that  the  elements  in  the  same  vertical  columns 
belong  to  a  group.  This  has  been  demonstrated  for  the  last  four 
columns  in  preceding  chapters;  it  will  be  proved  for  the  others 
later  on. 

In  the  light  of  these  facts,  we  are  led  to  the  conclusion  that 
the  physical  and  chemical  properties  of  the  elements  are  junctions  of 
their  atomic  weights;  and  when  we  consider  the  series  beginning 
with  lithium  and  sodium,  and  note  that  in  each  instance,  after  a 
difference  of  about  16,  there  follows  another  element  with  corre- 
sponding properties,  we  are  led  to  the  supposition  that  these 
properties  are  periodic  functions  of  the  atomic  weights. 

By  a  function  we  understand  in  general  a  dependent  relation  between 
two  or  more  magnitudes,  of  such  a  sort  that,  when  one  changes,  the 
other  does  likewise.  In  the  equations  y=a±x;  y=ax;  y=xn,  etc., 
y  is  a  function  of  x.  A  periodic  function  requires  that  the  same  value 
appear  for  one  magnitude  in  regular  intervals  as  the  other  magnitude 
steadily  increases.  An  example  of  this  kind  is  presented  by  the  gonio- 
metric  functions,  as  y=sm  x,  etc.,  for  every  time  x  increases  by  2?r,  y 
comes  to  have  the  same  value  again. 

If  we  desire  to  substantiate  the  conclusion  just  stated,  we  shall 
have  to  investigate  first  the  length  of  each  period,  in  other  words, 
determine  how  many  elements  intervene  in  the  table,  according  to 
increasing  atomic  weight  between  two  with  analogous  properties. 

It  has  already  been  shown  that  for  the  elements  as  far  as  chlo- 
rine, a  period  always  includes  seven  elements.  After  chlorine 
comes  potassium  (39),  which  thus  falls  into  the  column  under 
sodium.  The  following  elements, 

K39.2    Ca40.1     Sc44.1    T148.1    V51.2    O52.1    Mn55.0, 
correspond  very  well  with  the  preceding  series, 
Na23.1    Mg24.4    A127.1    Si28.4    P31.0    S32.1     C135.5- 


§213.]      THE  PERIODIC  SYSTEM    OF    THE   ELEMENTS.         299 

at  least  so  far  as  the  valence  and  the  form  of  the  compound  are 
concerned  (A1203  and  Sc2O3,  TiO2  and  SiO2,  K2CrO4  and  K2SO4, 
KMnCU  and  KCICU),  although  the  similarity  of  these  elements  in 
other  respects  is  not  very  marked. 

The  elements  following  manganese,  viz.,  Fe  55.9,  Co  59.0, 
Ni  58.7,  however,  do  not  fit  in  at  all  under  K,  Ca,  Sc;  but  if  we 
pass  these  by  there  follows  another  series  of  seven  elements,  which 
corresponds  to  the  one  beginning  with  potassium: 

Cu  63.6    Zn  65.4    Ga  70     Ge  72.5    As  75.0    Se  79.2    Br  80.0. 

We  therefore  reach  the  conclusion  that,  after  the  first  two  periods 
of  seven  elements  ending  with  chlorine  must  come  one  of  seventeen 
elements  (two  of  seven  each,  and  three  elements  placed  at  the  side), 
if  the  elements  in  the  same  vertical  column  are  to  correspond  in 
their  properties. 

This  large  period  of  seventeen  elements  can,  therefore,  be 
arranged  under  the  preceding  small  period  of  seven  elements  in 
the  following  way: 

SMALL   PERIOD. 

Na23.1  Mg24.4  A127.1    Si28.4  P31.0      S32.1    C135.5 

LARGE    PERIOD. 

K39.2  Ca40.1  Sc44.1  Ti48.1  V51.2  Cr52.1  Mn55.0  Fe55.9  Co59.0 
M58.7  Cu63.6  Zn65.4  Ga70  Ga72.5  As75.0  Se792  BrSO.O. 

In  order  to  arrange  in  periods  the  elements  whose  atomic  weights 
exceed  eighty,  it  is  again  necessary  to  assume  large  periods,  and, 
moreover,  to  leave  several  places  vacant.  In  this  manner  we 
arrive  at  the  scheme  known  as  MENDELEEFF'S  table  (see  p.  301). 

As  to  the  position  of  hydrogen  in  this  table  opinions  are  divided. 
MENDELEEFF  placed  this  element  in  the  first  group,  above  lithium;  its 
chemical  properties  indicate  without  doubt  that  it  belongs  with  these  j 
metals.  On  the  other  hand,  ORME  MASSON  has  presented  arguments 
for  placing  it  at  the  head  of  group  VII,  as  is  done  in  this  table.  These 
arguments  are  as  follows: 

(1)  The  molecule  of  hydrogen  contains  two  atoms,  as  does  a  halogen 
molecule,  while  the  molecule  of  an  alkali  metal  consists  of  one  atom. 
(2)  The  very  low  boiling-point  of  hydrogen  indicates  a  similarity  to 
the  halogens;  moreover  the  boiling-points  of  the  alkali  metals  fall  with 


300  INORGANIC   CHEMISTRY.  [§215. 

increasing  atomic  weight.  (3)  The  difference  between  the  atomic 
weights  of  the  elements  of  a  horizontal  series  is,  on  the  average,  3.  By 
placing  hydrogen  in  group  VII  it  differs  by  3  from  the  next  element, 
helium;  but  it  is  then  also  in  good  agreement  with  fluorine,  for  the 
mean  difference  in  atomic  weight  between  the  successive  elements  of  a 
column  is  16.  The  difference  F  -H  =18,  while  Li  -H  =6,  i.e.,  there  is  no 
analogy  in  the  latter  case.  (4)  Liquid  and  solid  hydrogen  have  no  metallic 
properties.  (5)  The  most  important  argument  for  placing  hydrogen 
in  the  first  group  is  based  on  its  relation  to  the  acids,  which  may  be 
regarded  as  salts  of  hydrogen.  But  in  organic  compounds  chlorine 
can  replace  hydrogen  without  essentially  altering  the  nature  of  the 
substance.  This  "organic"  argument  thus  offsets  the  ''inorganic" 
one,  based  on  the  analogy  of  acids  and  salts. 

As  to  the  elements  discovered  in  the  atmosphere  since  1894,  viz., 
helium  (at.  wt.  4),  neon  f20),  argon  (40  ,  krypton  (81.6)  and  xenon  (128), 
it  is  clear  that  they  form  a  natural  group,  for  their  properties  display 
great  analogy  (see  §  110).  Since  they  are  not  able  to  form  compounds 
with  other  elements,  they  can  be  regarded  as  nullivalent.  In  that  case 
their  group  could  find  a  place  after  the  eighth,  or  before  the  first,  group 
(compare  the  table,  page  294),  thus  forming  a  bridge,  or  transition, 
from  the  strongest  electro-negative,  to  the  strongest  electro-positive, 
elements.  However,  it  must  be  noted  that  argon  with  an  atomic  weight 
of  40  precedes  potassium  with  an  atomic  weight  of  39.  As  may  be  seen 
from  Plate  I,  their  atomic  volumes  fit  into  LOTHAB  MEYER'S  curve 
very  well. 

Group  VIII,  as  has  been  said,  owes  its  origin  to  the  setting 
at  the  side  of  the  elements  included  in  it,  for  by  this  means  the 
corresponding  elements  of  groups  I-VII  could  be  brought  under 
each  other.  It  will  thus  be  of  importance  to  the  system,  if  the 
nine  elements  of  this  group  display  so  much  analogy  to  each  other 
that  the  grouping  of  them  together  appears  actually  justified.  Now 
this  is  really  the  case,  as  is  seen  from  the  following  study  of  their 
properties : 

1.  All  these  elements  are  of  a  gray  color  and  difficultly  fusi- 
ble;   indeed,  osmium  is  one  of  the  hardest  of  all  metals  to  fuse 
(2500°);   iridium  melts  at  1950°,  wrought  iron  at  1500°,  etc.     The 
melting-point  of  iron  is  higher  than  that  of  cobalt,  and  the  latter 
higher  than  that  of  nickel.     A  similar  fall  of  this  constant  is  found 
with  ruthenium,  rhodium  and  palladium,  and  also  with  osmium, 
iridium  and  platinum. 

2.  Their  atomic  volumes  are  small  in  comparison  with  those 


§  213.]       THE  PERIODIC   SYSTEM    OF    THE   ELEMENTS.       301 


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302  INORGANIC    CHEMISTRY.  [§§  213- 

of  the  neighboring  elements.  The  atomic  volume  of  molybdenum 
is  11.2,  that  of  ruthenium,  rhodium  and  palladium  about  9;  that  of 
silver  10.3;  that  of  cadmium  13.0. 

3.  They  display  in  a  marked  degree  the  ability  to  let  hydrogen 
pass  through  at  red-heat,  or  to  condense  it  on  themselves  at  or- 
dinary temperatures.    The  former  property  is  especially  developed 
in  iron  and  platinum,  the  latter  in  palladium. 

4.  It  is  only  with  these  metals  that  we  find  RO4  compounds,  in 
other  words,  they  are  the  only  ones  which  can  be  octivalent.     The 
compounds  Os04  and  RuO4,  for  example,  are  known,  and  also  those 
with  carbon  monoxide,  as  Ni(CO)4  and  Fe(CO)4.     We  find  here 
a  general  tendency  to  form  compounds  with  carbon  monoxide, 
e.g.  PtCl2-3CO.     In  this  last  compound  platinum  can  also  be  con- 
sidered octivalent. 

5.  They  form    stable  and,   in    most    cases,  well-crystallizing 
double    salts    with    potassium    cyanide.     Iron,    ruthenium,    and 
osmium  give  compounds  of  the  type  K4R(CN)6;  cobalt,  rhodium, 
and  iridium  form  K3R(CN)6,  while  the  elements  of  the  last  column, 
nickel,  palladium,  and  platinum,  give  K2R(CN)4  double  salts. 

6.  They  all  form  colored  salts:  those  of  cobalt  are  red  or  blue; 
the  nickel  salts  are  green;   all  the  rest  are  of  various  shades  of 
brown. 

7.  They  all  possess   the   ability   to    condense   on   themselves 
other  gases  than  hydrogen  in  larger  or  smaller  amounts ;  especially 
is  this  true  of  the  platinum  metals.     Platinum  and  palladium 
absorb  carbon  monoxide  greedily. 

Let  us  now  examine  the  seven  other  groups  (vertical  col- 
umns) more  closely.  If  we  bear  in  mind  what  was  stated  with 
reference  to  the  large  periods,  it  is  apparent  that  not  all  the  ele- 
ments in  such  a  group  display  perfect  chemical  analogy.  Such 
analogy  is  found,  however,  when  we  compare  with  each  other 
only  those  elements  that  belong  to  the  even  or  the  odd  rows  of 
the  large  periods.  The  similarity  of  the  elements  of  these  divisions 
is  seen,  in  the  case  of  the  large  periods,  from  the  following  facts: 

1.  Only  the  elements  of  the  odd  rows  give  hydrogen  or  alkyl 
compounds. 

2.  In  the  even  rows  the  basic  properties  of  the  hydroxides  are 
prominent,  in  the  odd  rows  the  acidic  properties. 

In  general  it  may  be  said  that,  passing  from  left  to  right  in  the 


215.]          THE   PERIODIC   SYSTEM   OF    THE   ELEMENTS.      303 

table,  we  meet  first  those  that  form  the  strongest  bases  and  then 
gradually  those  that  give  acids.  The  latter  property  is  most 
marked  in  the  halogens,  since  they  even  form  strong  acids  on 
combination  with  hydrogen. 

A  similar  change  is  observed  in  going  from  the  top  to  the  bottom 
of  the  system.  As  the  atomic  weight  increases,  the  metallic  (base- 
forming)  nature  of  the  elements  in  each  group  becomes  more 
predominant. 

214.  The  periodic  system  of  the  elements  is  of  value  in  four 
different  respects: 

1.  In  constructing  a  system  of  the  elements. 

2.  In  ascertaining  the  atomic  weights  of  elements  whose  equiva- 
lent weights  only  could  be  determined. 

3.  In  foretelling  the  properties  of  elements  as  yet  undiscovered. 

4.  In  confirming  or  correcting  atomic  weights. 

Let  us  look  at  these  various  applications  in  some  detail: 

The  Use  of  the  Periodic  Law  in  Constructing  a  System  of  the 

Elements. 

215.  After  a  careful  comparison  of  the  elements  MENDELEEFF 
reached  the  following  important  conclusion:   The  entire  character  of 
an  element,  as  displayed  in  its  physical  as  well  as  in  its  chemical 
properties,  is  determined  by  the  position  which  it  occupies  in  the  sys- 
tem and  particularly  by  the  four  adjacent  elements,  the  atomic 
analogues.     If  an  element  is  in  an  even  series,  the  elements  in 
the  adjoining  even  series  are  its  atomic  analogues;  the  same  is  true 
of  the  odd  rows.     From  this  it  follows  that,  when  the  properties  of 
an  element  are  exactly  known,   its  place  in  the  system  can  be 
assigned.    A  couple  of  illustrations  will  make  this  clear. 

The  element  beryllium  possesses  marked  analogy  to  alumin- 
ium on  the  one  hand  and  magnesium  on  the  other;  therefore  it 
was  a  much  discussed  question  whether  its  oxide  should  be  given 
the  formula  BeO  or  ~Be2Oz-  Since,  according  to  analysis,  9.1 
parts  Be  combine  with  16  parts  O,  the  atomic  weight  would  be 
in  the  former  case  9.1,  in  the  latter  |X  9. 1  =  13. 7  With  the  atomic 
weight  13.7  the  element  would  stand  between  nitrogen  and  oxygen; 
as  nitrogen  and  the  elements  of  the  sulphur  group  as  well  yield 
only  acid-forming  oxides,  beryllium  oxide  would  have  to  be  an 
acid  anhydride  too,  which  is  not  the  case,  it  being  a  basic  oxide 


304  INORGANIC    CHEMISTRY.  [§§  215- 

Thus  beryllium  would  not  fit  in  the  system  with  that  atomic  weight. 
If,  however,  it  has  the  atomic  weight  9.4,  it  falls  in  the  horizontal 
series  between  the  strongly  basic  lithium  and  the  weakly  acidic 
boron  and  over  magnesium,  with  which  it  shows  real  analogy. 
This  is  indeed  its  fit  place,  i.e.  its  properties  are  those  which  must 
belong  to  an  element  in  this  position  (see  table).  Farther,  the 
properties  of  beryllium  are  to  those  of  lithium  as  the  properties 
of  boron  are  to  those  of  beryllium,  or,  in  the  form  of  a  proportion: 
Be: Li:  :B:Be.  Just  as  lithium  forms  a  more  strongly  basic  oxide 
than  beryllium,  so  the  basic  character  of  beryllium  oxide  is  stronger 
than  that  of  boron  oxide;  again,  beryllium  chloride  is  more  vola- 
tile than  lithium  chloride,  boron  chloride  more  volatile  than 
beryllium  chloride. 

We  also  have  the  relation  Be:Mg:  :Li:Na:  :B:A1,  for  beryllium 
oxide  is  less  basic  than  magnesium  oxide,  lithium  oxide  than 
sodium  oxide,  boron  oxide  than  aluminium  oxide.  Beryllium 
fluoride  dissolves  in  water,  magnesium  fluoride  does  not;  simi- 
larly boron  fluoride,  but  not  aluminium  fluoride,  is  soluble  in  water. 

Finally  we  have  Be:Al:  :Li:Mg:  :B:Si.  The  hydroxides  of 
beryllium  and  aluminium  are  very  similar  to  each  other;  they  are 
gelatinous  and  soluble  in  alkalies.  Both  metals  are  scarcely  acted 
upon  by  nitric  acid  and  both  dissolve  in  alkalies  with  the  evolution 
of  hydrogen.  Their  chlorides  must  be  prepared  in  the  same  way 
from  the  oxides  (by  heating  with  charcoal  in  a  current  of  chlorine). 
Likewise  lithium  and  magnesium  are  analogous  in  certain  respects; 
the  carbonate  and  the  phosphate  of  lithium,  like  the  corresponding 
salts  of  magnesium,  are  very  difficultly  soluble,  which  is  in  marked 
contrast  with  the  other  metals  of  the  lithium  group.  Boron  and 
silicon  both  form  very  refractory  oxides  and  salts;  their  fluorine 
compounds  are  decomposed  by  water  in  a  similar  manner,  etc. 

The  evidence  in  accordance  with  which  beryllium  was  assigned  its 
present  position  in  the  system  was  subsequently  confirmed  directly 
by  the  determination  of  the  vapor  density  of  the  chloride,  which  led 
to  the  formula  BeCl2.  The  vapor  density  of  the  beryllium  compound 
of  an  organic  substance  (acety  lace  tone)  also  led  to  9.1  as  the  atomic 
weight  of  beryllium. 

As  a  second  example  let  us  take  thallium.  This  element  dis- 
plays analogy  with  the  alkali  metals  and  also  with  aluminium,  as 


216.]        THE  PERIODIC  SYSTEM  OF   THE  ELEMENTS.  305 

well  as  with  lead  and  mercury.  According  to  its  atomic  weight  it 
must  lie  between  the  two  latter  and  belongs  in  the  aluminium  group. 
This  position  is  justified  when  the  following  relation  is  taken  into 
consideration : 

Tl:Al::Hg:Mg::Pb:Si. 

The  highest  stages  of  oxidation  show  unmistakable  analogy, 
since  the  oxides  have  the  same  properties  by  pairs.  The  oxides  of 
aluminium  and  thallium  are  weakly  basic,  those  of  mercury  (ic) 
and  magnesium  more  strongly  basic,  while  lead  dioxide  and  silica 
are  slightly  acidic.  Thallium,  mercury  and  lead  all  form  lower 
oxides  of  strongly  basic  character  (T^O,  Hg2O,  PbO,)  which  alumin- 
ium, magnesium  and  silicon  are  unable  to  do.  The  oxide  T^O 
is  comparable  with  l^O  in  its  properties,  the  lead  monoxide  with 
calcium  oxide  (T1:K:  :Pb:Ca);  see  the  table.  In  regard  to  the 
physical  properties,  it  should  be  observed  that  in  the  matter  of 
volatility  thallium  lies  between  mercury  and  lead. 

Use  of  the  System  in  Ascertaining  the  Atomic  Weights  of  Ele- 
ments whose  Equivalent  Weights  only  are  known. 

216.  A  good  instance  is  that  of  the  very  rare  element  indium. 
When  the  periodic  system  was  established  only  the  equivalent 
weight  of  indium  was  known,  the  analysis  of  the  chloride  having 
shown  it  to  be  38.3,  i.e.  it  was  known  that  38.3  parts  of  indium  com- 
bine with  35.5  parts  of  chlorine.  If  this  were  the  atomic  weight 
and  consequently  In2O  the  formula  of  the  oxide,  the  element  would 
have  to  occupy  the  place  which  potassium  now  has.  Not  only  is 
this  place  taken,  but  the  oxide  is  only  weakly  basic,  while  it  should 
have  strongly  basic  properties  if  it  belonged  to  the  first  group.  If 
the  atomic  weight  were  76.6,  corresponding  to  the  oxide  InO,  the 
element  would  stan  T  between  arsenic  and  selenium  in  the  table;  its 
oxide  could  not  then  have  the  formula  InO,  but  would  have  to  be 
In2Os  or  In2Os  and  have  acid  properties.  On  the  other  hand,  there 
is  no  place  for  a  metal  with  the  atomic  weight  76.6  and  an  oxide  InO 
hi  the  second  group,  where  the  oxides  have  this  type.  If  the  oxide 
be  In2Gs,  the  atomic  weight  of  indium  must  be  115.  Now,  as  a 
matter  of  fact,  there  is  a  vacant  place  in  the  system  between 
Cd  =  112  and  Sn=119  for  an  element  with  an  I^Os  oxide.  In 


306  INORGANIC    CHEMISTRY.  [§§  216- 

the  same  way  as  we  have  shown  it  for  the  oxides  R2O  and  HO,  it  can 
also  be  demonstrated  that  indium  could  not  be  located  in  the  table 
with  an  oxide  InC>2,  In2O5,  InOs,  etc.  There  remains,  thus,  no 
other  possibility  than  to  give  the  oxide  the  formula  In2O3  and  the 
element  an  atomic  weight  of  115. 

Let  us  now  see  whether  the  properties  of  the  element  and  its 
compounds  conform  to  this  location.  Two  of  its  atomic  analogues 
for  this  position  are  cadmium  and  tin.  The  oxides  of  both  are  easily 
reduced;  this  is  true  also  of  indium  oxide.  If,  now,  we  consider 
the  metals : 

Ag,  Cd,  In,  Sn,  Sb  (7th  series), 

we  notice  first  that  the  melting-point  of  silver  lies  higher  than  that 
of  cadmium;  likewise  antimony  melts  higher  than  tin: 

Ag>Cd;  Sn<Sb. 

If  indium  fits  in  this  series,  it  must,  therefore,  have  the  lowest 
melting-point  of  these  five  metals.  This  is  actually  the  case;  its 
melting-point  is  176°.  In  the  color  of  the  metals  there  is  a  further 
analogy;  silver,  cadmium  and  tin  are  white,  and  so  is  indium.  As 
to  the  specific  gravity,  cadmium  has  the  specific  gravity  8.6,  silver 
10.5;  the  difference  is  thus  1.9.  The  difference  of  the  specific 
gravities  of  tin  and  antimony  is,  on  the  contrary,  small, 
7.3  —  6.7=0.6.  If  indium  stands  between  these  two  pairs  of 
elements  its  specific  gravity  should  be  smaller  than  the  mean  be- 
tween cadmium  and  tin,  i.e.  J(8.6  +  7.3)  =7.9.  In  reality  its 
specific  gravity  is  7.42. 

The  position  of  the  metal  requires  that  its  oxide  be  feebly  basic, 
weaker  even  than  the  oxides  of  cadmium  and  thallium  (T^Os), 
since  cadmium  is  at  the  left  and  thallium  farther  down  in  the  system 
(§  215) ;  on  the  other  hand,  it  must  be  more  strongly  basic  than 
the  oxides  of  aluminium  and  tin.  This  condition  also  is  fulfilled, 
as  the  following  sets  forth:  the  oxides  of  aluminium  and  tin,  be- 
cause of  their  slightly  acid  nature,  dissolve  in  alkalies,  forming  defi- 
nite compounds  with  the  latter.  The  oxides  of  cadmium  and 
thallium,  however,  are  insoluble  in  alkalies;  they  are  distinctly 
basic.  Indium  seequioxide  is  soluble  in  alkalies,  but  forms  no 
definite  compound  with  them. 

Finally,  there  is  also  the  conduct  of  indium  salts  toward  hydro- 
gen sulphide  which  supports  the  placing  of  the  element  between 


217.]        THE  PERIODIC   STSTEM   OF   THE  ELEMENTS.         307 

cadmium  and  tin,  for  indium  also  is  precipitated  by  hydrogen 
sulphide  from  acid  solutions. 

Prediction  of  the  Properties  of  Elements  yet  Undiscovered. 

217.  When  MENDELEEFF  constructed  his  table,  the  elements 
gallium  and  germanium  (fifth  series)  were  still  unknown.  From 
the  properties  of  the  atomic  analogues  he  ventured  however  at  that 
time  to  predict  the  properties  of  these  elements;  thirteen  years 
later  WINKLER  discovered  an  element  with  the  atomic  weight  72 
(MENDELEEFF  had  termed  such  an  one  "  ekasilicon";  the  discoverer 
named  it  germanium).  Indeed  there  proved  to  be  a  very  close 
agreement  between  the  actual  properties  and  those  predicted,  as 
the  following  summary  evidences: 


Properties  of  ekasilicon  predicted  by 

MENDELEEFF. 


Properties  of  germanium  discovered  by 
WINKLER. 


1.  The  atomic  weight  must  be  the 

mean  of  the  four  atomic  ana- 
logues Si,  Sn,  Zn,  Se,  i.e_ 
K28.4  +  118.5+65.4  +  79.1)  =72.85. 

2.  The  specific  gravity  (deduced  as 

for  indium  above)  must  be  5.5. 

3.  The  atomic  volume  must  lie  be- 

tween those  of  silicon  (13)  and 
tin  (16),  but  be  only  a  trifle 
above  13. 

4.  Since  it  belongs  in  an  odd  series,  it 

must  be  able  to  form  alkyl  com- 
pounds. Judging  from  analogies, 
the  boiling-point  of  Es(C2H5)4 
must  be  160°,  its  specific  gravity 
0.96. 

5.  The  acid  properties  of  EsO?  must 

be  stronger  than  those  of  SnO2. 

6.  The  specific  gravity  of  EsO2  is  4. 7. 

7.  Since  the  oxides  of  indium  and 

arsenic  are  easily  reduced,  this 
must  also  be  true  of  EsO2. 

8.  EsS2,  because  of  its  analogy  with 

SnS9.  will  probably  be  soluble  in 
NH4SH. 

9.  EsCl4  is  liquid,,  boils  below  100° 

(since  the  boiling-point  of  SiCl4  is 
57°  and  that  of  SnCl4  115°),  and 
has  a  specific  gravity  of  1.9. 
10.  K2SnF6  being  more  readily  soluble 
in  water  than  K2SiF0,  the  solu- 
bility of  K2EsF6  must  also  be 
greater  than  that  of  K2SiFe. 


1.  At.  wt- 72.5. 

2.  Sp.  g.=  5.469  at  20°. 

3.  At.  vol.=13.a 


4.  Alkyl  compounds  were  obtained. 
Ge(C2H5)4  boils  at  160°  and  its 
specific  gravity  is  a  little  less 
than  1. 


5.  GeO2  lacks  entirely  the  basic  prop- 

erties which  are  found  to  a  lim- 
ited extent  in  SnO2. 

6.  The  specific  gravity  of  GeO2  is 

4.703  at  18°. 

7.  GeO2  is  easily  reduced  to  the  metaj 

by  heating  with  carbon  or  in 
hydrogen. 

8.  GeS2  dissolves  readily  in  NH4SH. 


9.  GeCl6  is  liquid,  boils  at  86°,  and 
has  a  specific  gravity  of  1.887. 


10.  KjSiF6  is  almost  insoluble- 
K2GeF8  dissolves  hi  34  parts  of 
boiling  water. 


308  INORGANIC    CHEMISTRY.  [§§  218- 

Use  of  the  Periodic  Syst3m  in  Correcting  Atomic  Weights. 

218.  In  the  group  of  the  platinum  metals  the  atomic  weights 
were  previously  determined  by  BERZELIUS  and  FREMY    to  be  as 
follows:   Os  =  199,  Pt  198,  Ir  197.     In  this  order  the  metals  named 
do  not;  however,  fit  into  the  system,  for   osmium  should  be  the 
first  of  the  three  on  account  of  its  analogy  with  iron  and  ruthenium. 
i.e.  it  should  have  the  smallest  atomic  weight  of  the  three.      On 
the  other  hand,  platinum,  which  is  more  akin  to  palladium,  ought 
to  have  the  highest  atomic  weight.     A  painstaking  investigation 
by  SEUBERT  showed  that  in  reality  osmium  has  the  atomic  weight 
191,  iridium  193.0,  and  platinum  194.8,  which  order  is  in  harmony 
vith  the  system. 

219.  Graphic  representation. — The  fact  that  an  arrangement 
of  the  elements  according  to  increasing  atomic  weight  also  makes 
the  gradual  change  of  the  physical    properties  apparent  can  be 
seen  most  clearly  from  a  graphic  representation  (as  proposed  by 
LOTHAR  MEYER),  in  which  the  atomic  weights  are  the  abscissas 
and  the  atomic  volumes  the  ordinates  (see  Plate  I  at  end  of  book). 

The  first  thing  we  notice  in  the  curve  is  the  regular  rise  and 
fall  of  atomic  volumes.  At  the  beginning  of  each  period  the 
atomic  volume  is  at  a  maximum;  it  reaches  a  minimum  half  way 
through  the  period  (in  the  large  periods  at  group  VIII)  and  then 
increases  again.  In  the  descending  portions  are  the  ductile,  on 
the  ascending  the  brittle,  elements.  On  the  ascending  portions 
and  at  the  maxima  are,  further,  the  gaseous  and.  the  easily  fusible 
elements;  on  the  descending  and  at  the  minima  the  difficultly 
and  very  difficultly  fusible.  On  the  descending  portions  are  the 
electro-positive,  on  the  ascending  the  electro-negative,  elements. 
The  periodicity  of  the  elements  thus  becomes  very  evident. 

220.  The  discovery  of  the  periodic  system  once  more  thrust  into  the 
foreground  of  interest  one  of  the  oldest  of  questions,  viz.,  that  concerning 
the  unity  of  matter. 

The  striking  connection  between  all  the  properties  of  the  elements  and 
their  atomic  weights  leads  unavoidably  to  the  assumption  of  a  primordial 
or  ground  substance.  As  to  the  nature  of  this  substance  no  evidence  is  at 
present  obtainable.  PROUT,  in  1815,  regarded  hydrogen  as  such  a  sub- 
stance. He  observed  that  the  (then  accepted)  atomic  weights  of  many 
elements,  based  on  hydrogen  as  unity,  are  whole  numbers.  Later,  very 
accurate  atomic  weight  determinations,  particularly  those  of  STAS,  which 


221.]          THE    PERIODIC   SYSTEM    OF    THE    ELEMENTS.      309 

were  undertaken  with  the  purpose  of  testing  this  hypothesis,  demonstrated 
with  certainty,  however,  that  it  was  untenable.  Cf.  §  267. 

221.  The  periodic  system  of  the  elements  is  one  of  the  most  important 
discoveries  in  the  field  of  inorganic  chemistry;  it  can  never  lose  its  impor- 
tance, though  it  is  gradually  becoming  more  evident  that  the  system 
in  its  present  form  represents  the  relations  of  the  elements  to  each  other 
merely  in  an  approximate  way  and  is  only  a  crude  first  attempt  at  a  real 
system. 

There  are  indeed  serious  objections  to  the  periodic  sys- 
tem. These  objections  concern,  in  the  first  place,  the  positions  which 
certain  elements  occupy  in  the  system,  and  which  agree  very  poorly  with 
their  properties.  This  is  the  case,  e.g.,  with  gold  and  copper,  which,  indeed, 
show  some  analogy  with  lithium  and  sodium  in  their  ows-compounds,  but 
otherwise  differ  decidedly  from  the  latter  elements.  The  same  is  true  of  the 
metals  of  the  cerium  group  (Ce,  La,  Nd,  Pr,  etc.).  The  analogy  with  the  other 
elements  of  the  same  group  is  feeble.  If  the  reply  be  made  that  these  elements 
are  still  too  little  known  because  of  their  rarity  and  the  great  expenditure  of 
energy  and  pains  required  for  their  investigation,  it  must  be  protested  that 
the  same  is  also  true  of  those  better  known,  such  as  cerium,  and  that  if  these 
elements  should  disclose  themselves  as  complexes  of  several  (which  is  not 
improbable),  it  is  very  doubtful  whether  there  would  be  a  place  in  the  system 
for  elements  with  approximately  those  properties  which  we  are  at  present 
acquainted  with  in  lanthanum,  neodymium,  and  praseodymium,  unless  more 
than  one  element  were  put  in  a  place,  as  has  been  proposed. 

A  second  objection  concerns  the  inability  to  fit  all  elements  into  the 
system.  This  is  particularly  the  case  with  tellurium.  Its  physical  and 
chemical  properties  put  it  without  doubt  in  the  sulphur  group,  and  here 
there  is  a  space  for  it,  if  its  atomic  weight  were  only  about  125,  or  at  least 
smaller  than  that  of  iodine  (126.92).  Nevertheless  repeated  and  careful 
investigations  have  fixed  its  atomic  weight  at  127.5.  The  same  difficulty 
presents  itself  with  cobalt  and  nickel.  According  to  their  atomic  weights 
the  four  elements  Fe,  Co,  Ni,  Cu  must  be  arranged  as  follows:  Fe,  55.85, 
Ni  58.68,  Co  58.97,  Cu  63.57.  But  the  order  which  best  corresponds  with 
their  properties  is  the  one  given  first,  cobalt  belonging  more  strictly  with 
iron,  nickel  with  copper. 

When  we  recall  that  the  newly  discovered  elements  of  the  argon  group 
fit  into  the  system  very  satisfactorily  and  that  radium,  too,  finds  a  place  in 
it,  we  have  little  reason  to  believe  that  the  usefulness  of  the  Periodic  System 
is  in  any  way  exhausted. 


310  INORGANIC    CHEMISTRY.  [§§  222 

LITHIUM  AND  SODIUM. 

Lithium. 

222.  This  metal  is  not  found  free  in  nature;  in  combination,  however, 
it  is  very  widely  distributed,  although  always  in  small  amounts.  Many 
mineral  waters  contain  it.  It  occurs  chiefly  in  the  silicate  lepidolite,  or 
lithia-mica,  also  as  the  phosphate  in  triphylite  and  in  company  with  alu- 
minium, sodium,  and  fluorine  in  amblygonite.  Finally,  lithium  is  met 
with  in  the  ashes  of  certain  plants,  such  as  tobacco,  indicating  that  it  is 
also  contained  in  the  soil.  With  the  aid  of  the  spectroscope  it  can  be 
detected  in  very  many  minerals. 

Lithium  can  be  obtained  from  lepidolite  by  the  following  very  simple 
process :  The  mineral  is  fused  and  then  poured  into  cold  water,  whereupon 
it  becomes  very  brittle  and  its  silicates  are  brought  into  such  a  condition 
that  they  can  be  decomposed  by  hydrochloric  acid.  The  finely  powdered 
mass  is  boiled  with  hydrochloric  acid  and  the  metals  Ca,  Al,  Mg,  etc., 
precipitated  by  soda  from  the  resulting  solution  (after  filtering  off  the 
silica),  lithium  and  the  other  alkali  metals  remaining  in  solution.  By 
evaporation  a  salt  mixture  is  obtained  from  which  the  lithium  chloride 
can  be  isolated  by  extraction  with  alcohol,  the  insoluble  chlorides  of 
sodium  and  potassium  remaining  behind. 

Metallic  lithium  is  prepared  by  electrolysis  of  the  fused  chloride  or  a 
concentrated  solution  of  this  salt  in  pyridine.  Next  to  solid  hydrogen, 
it  is  the  lightest  of  all  solid  substances,  its  specific  gravity  being  only 
0.59,  so  that  it  floats  on  coal-oil,  It  is  silvery-white,  but  tarnishes  very 
rapidly  in  moist  air.  Melting-point  180°.  When  heated  in  the  air  it 
burns  with  an  intense  white  light  to  the  oxide;  at  ordinary  temperatures 
it  is  not  so  readily  oxidized  as  sodium  and  potassium.  It  decomposes 
water  with  the  evolution  of  hydrogen-,  the  heat  generated  is  sufficient, 
however,  to  melt  the  metal. 

Lithium  oxide,  Li20,  and  hydroxide,  LiOH.  The  former  is  obtained 
by  heating  the  nitrate  strongly.  It  dissolves  in  water  slowly,  forming 
the  hydroxide.  The  latter  is  a  white,  crystalline  substance,  similar 
to  caustic  soda;  it  dissolves  in  water,  producing  a  strongly  alkaline 
solution. 

Lithium  chloride,  LiCl,  crystallizes  anhydrous  in  regular  octahedra; 
below  0°  it  takes  up  two  molecules  of  water  of  crystallization  however. 
It  dissolves  very  easily  in  water  and  deliquesces  in  moist  air. 

Lithium  carbonate,  Li2CO3,  unlike  the  carbonates  of  the  other  alkalies, 
is  difficultly  soluble  in  water  (100  parts  of  water  at  13°  take  up  0.769 
part) ;  hence  it  can  be  precipitated  from  the  concentrated  solution  of 
the  chloride  by  ammonium  carbonate. 


223.]  LITHIUM   AND   SODIUM.  311 

Lithium  phosphate,  Li3PO4,  is  likewise  very  sparingly  soluble  in  water 
(1  part  in  2539  parts  of  water),  although  the  phosphates  of  the  other 
alkalies  are  freely  soluble.  The  formation  of  this  salt  serves  as  a  test  for 
lithium. 

The  lithium  spectrum  consist?  of  two  red  bands,  one  of  wnich  in  par- 
ticular is  easy  to  recognize. 

Sodium. 

223.  Sodium  occurs  in  nature  in  enormous  quantities1  and  is 
very  widely  diffused.  It  is  a  constituent  of  countless  silicates 
and,  as  a  result  of  rock  decay,  gets  into  the  soil,  whence  it  enters 
the  plants  and  finally  reaches  the  animal  organism.  The  nitrate 
is  known  as  Chili  saltpetre,  the  chloride  as  rock-salt  or  halite,  the 
carbonate  as  soda;  the  cryolite  (ice-stone)  of  Greenland  is  a  sodium 
aluminium  fluoride.  Common  salt,  Nad,  constitutes  the  main 
part  of  the  saline  matter  in  sea-water.  Certain  bodies  of  water 
such  as  the  Dead  Sea  of  Palestine,  and  the  Great  Salt  Lake  in 
North  America,  are  almost  saturated  solutions  of  common  salt. 

The  metal  was  first  obtained  by  DAVY  in  1807  by  the  elec- 
trolysis of  molten  sodium  hydroxide.  GAY-LUSSAC  and  THENARD 
got  it  by  heating  sodium  hydroxide  with  powdered  iron  to  white 
heat.  The  first  named  method  is  the  one  now  generally  employed 
for  its  commercial  manufacture,  inasmuch  as  electric  power  can 
be  obtained  quite  cheaply. 

For  this  purpose  sodium  hydroxide  is  heated  a  little  above  its  melting- 
point.  The  sodium  formed  at  the  cathode  is  kept  away  from  the  anode 
by  an  iron  net.  At  the  anode  hydroxyl  groups  are  liberated,  which  yield 
water  and  oxygen.  The  latter  escapes  but  the  water  dissolves  in 'the  molten 
mass  and  comes  in  contact  with  the  sodium  at  the  cathode,  causing  half 
of  it  to  be  changed  back  to  sodium  hydroxide,  while  hydrogen  js  -evolved. 
As  a  result  the  maximum  yield  of  metal  for  a  given  quantity  of  electricity 
is  only  50%.  If  the  temperature  gets  too  high  during  the  electrolysis, 
sodium  dissolves  in  the  molten  mass  and  is  oxidized  at  the  anode. 

Sodium  is  silvery-white,  melts  at  95.6°  and  boils  at  900°, 
turning  at  the  latter  temperature  to  a  colorless  vapor.  At 
ordinary  temperatures  it  is  very  soft,  so  that  it  can  be  readily 
cut  with  a  knife.  It  can  also  be  easily  pressed  through  a  small 
hole,  coming  out  in  the  form  of  wire.  Sp.  g.  13.5°  at  =0.9735. 

Sodium,  like  the  other  alkali  metals  (Li,  K,  Rb,  Cs),  dissolves  in  liquid 
ammonia.  If  one  of  these  metals  is  introduced  into  liquid  ammonia,  the 
bright  surface  of  the  metal  becomes  tarnished  with  an  indigo-blue  color, 
which  soon  turns  to  a  pretty  metallic  red.  The  metal  then  liquefies  and 
forms  a  bronze-colored  solution,  which  is  deep  blue  at  greater  dilution.  If 


312  INORGANIC   CHEMISTRY.  >  223- 

the  pressure  of  the  ammonia  over  the  solution  is  diminished  or  the  tem- 
perature raised,  the  solution  gives  off  ammonia  and  deposits  copper-red, 
crystalline  masses.  When  there  is  no  longer  a  liquid  solution  these  masses 
also  lose  ammonia  and  the  metal  is  left  behind  in  the  crystalline  form. 

RUFF  has  demonstrated  that  we  do  not  have  here  a  case  of  compounds 
being  formed  between  the  metal  and  the  ammonia,  but  that  the  copper-red 
masses  are  mixtures  of  metal  and  saturated  liquid  solution;  for  the  solution 
can  be  pressed  out,  leaving  a  compact  piece  of  metal. 

The  molecule  of  sodium  contains  only  one  atom,  as  is  proved 
by  the  depression  of  the  freezing-point  of  its  solution  in  tin.  A 
great  many  metals  have  this  same  property. 

In  moist  air  the  bright  surface  of  a  freshly  cut  piece  tarnishes 
rapidly,  but  in  air  that  has  been  dried  with  phosphorus  pentoxide 
it  keeps  its  metallic  lustre  for  days.  Sodium  can  be  heated  in  the 
air  to  melting  and  even  still  higher  without  catching  fire.  It 
ignites  only  when  heated  strongly,  whereupon  it  burns  with  a  very 
bright  yellow  light  (especially  in  an  atmosphere  of  oxygen).  With 
water  it  generates  hydrogen,  sodium  hydroxide  being  also  formed. 
If  it  is  held  firmly  in  one  place  during  this  process  (e.g.  by  laying 
it  on  a  piece  of  filter-paper  floating  on  water,  or  upon  ice),  the 
hydrogen  takes  fire  because  of  the  localization  of  the  heat  produc- 
tion. 

Sodium  finds  extensive  use  in  the  laboratory  and  in  the  arts. 
Because  of  its  strong  reduc ing-power  it  is  often  used  to  obtain 
the  elements  from  their  oxides;  magnesium  and  aluminium  were 
formerly  obtained  thus.  In  organic  chemistry  it  is  also  frequently 
employed  for  various  purposes. 

OXIDES  AND  HYDROXIDES  OF  SODIUM. 

224.  On  burning  sodium  in  dry  oxygen  a  mixture  of  two  oxides, 
Na^O  and  Na2O2,  results.  Sodium  oxide,  Na2O,  is  obtained 
pure  by  the  partial  and  slow  oxidation  of  sodium  with  oxygen 
under  reduced  pressure  and  removal  of  the  excess  of  metal  by 
distillation  in  a  vacuum.  The  oxide  dissolves  slightly  in  the 
metal  and  after  distilling  off  the  latter  the  oxide  is  left  in  the 
crystalline  form.  It  is  white;  it  dissolves  in  water,  forming 
sodium  hydroxide,  NaOH,  and  giving  off  much  heat. 

The  peroxide,  Na202,  is  obtained  by  heating  sodium  in  a 
current  of  oxygen  till  no  more  oxygen  is  absorbed.  With  8  mols. 


224.]  OXIDES   AND  HYDROXIDES  OF  SODIUM.  313 

water  it  forms  a  hydrate,  Na202  +  8H2O.  Since  it  yields  hydrogen 
peroxide  with  dilute  acids  and  is  a  vigorous  oxidizing-agent  it  is 
manufactured  commercially. 

Sodium  hydroxide,  NaOH;  caustic  soda,  is  formed,  together 
with  metallic  sodium,  when  sodium  monoxide  is  reduced  in  a  cur- 
rent of  hydrogen.  The  ordinary  method  of  preparing  caustic  soda 
consists  in  boiling  soda  with  slaked  lime: 

Na2CO3  +  Ca(OH)2  =2NaOH+CaCO3; 
or  2Na-  +  CO3" + Ca"  +  20H'  =  2(Na-  +  OH')  +  CaCO3. 

•As  the  solubility  product  (§  73)  of  CaC03  molecules  is  very  small, 
the  ions  Ca"  and  C03"  must  unite  and  the  carbonate  of  lime  sinks 
to  the  bottom.  In  order  to  make  the  decomposition  of  sodium 
carbonate  complete,  a  slight  excess  of  slaked  lime  is  added.  Never- 
theless, the  solution  does  not  contain  an  appreciable  quantity  of 
calcium  hydroxide.  The  reason  of  this  is  clear:  In  the  solution 
there  are  a  large  number  of  OH-ions;  as  a  result  the  number 
of  Ca-ions  can  only  be  very  small,  for  the  value  of  the  solubility 
product  of  calcium  hydroxide  is  reached  with  even  a  very  low 
concentration  of  the  latter  ions. 

Sodium  hydroxide  is  now  manufactured  on  a  large  scale  by  the 
electrolysis  of  concentrated  brine.  This  method  yields  an  almost 
chemically  pure  hydroxide  and  it  dominates  the  market  with  users 
of  high-grade  caustic.  Three  types  of  electrolytic  processes  are  in 
operation:  the  diaphragm  type,  the  amalgamation  type,  and  the 
bell  type.  In  the  first  type  the  cathode  and  anode  compartments 
are  separated  by  a  diaphragm.  In  the  Griesheim  process,  a  suc- 
cessful representative  of  this  type,  the  cathode  is  of  iron  and  the 
anode  ferrous-ferric  oxide,  Fe304>  that  has  been  fused  at  2000°- 
3000°  and  cast  into  plates.  (This  makes  an  anode  unaffected  by 
chlorine.)  In  a  freshly  charged  bath  containing  only  chloride  solu- 
tion the  current  is  carried  mainly  by  sodium  and  chlorine  ions; 
but  as  fast  as  the  sodium  is  liberated  at  the  cathode  and  reacts 
with  water  to  form  hydroxide  and  free  hydrogen,  the  ions  of 
the  hydroxide  participate  in  the  transport  of  the  current. 
The  sodium  atoms  are  reliberated  and  again  react  with  water  to 
form  hydrogen,  while  the  discharged  hydroxyl  ions  yield  water 
and  oxygen.  As  a  net  result  of  the  electrolysis  we  have,  so  to 
speak,  an  intentional  decomposition  of  alkali  chloride  accom- 


314  INORGANIC   CHEMISTRY.  [§§  224- 

panied  in  an  increasing  measure  by  an  unintended  decomposition 
of  water.  On  this  account  the  chloride  electrolysis  cannot  be 
carried  to  completion;  in  practice  the  process  is  interrupted 
as  soon  as  the  alkali  hydroxide  concentration  gets  up  to  8%. 
The  caustic  cathode  liquid  is  then  replaced  by  fresh  brine  and  the 
former  is  evaporated  in  vacuum  pans  to  a  concentration  of  50%, 
whereupon  the  undecomposed  chloride  separates  out  and  is 
returned  to  the  electrolytic  cell. 

The  diaphragm  process  most  favorably  known  in  America  is 
the  TOWNSEND  process.  The  compartment  containing  the 
(graphite)  anodes  occupies  the  center  of  the  cell  and  on  each  side 
i:  a  diaphragm  of  tliin  asbestos.  The  cathodes,  of  woven  wire, 
:est  closely  against  the  outside  of  the  diaphragm  and  the  cathode 
compartment  is  filled  with  warm  oil,  that  is  kept  in  lively  cir- 
culation by  the  escaping  hydrogen.  The  freshly  formed  caustic 
liquor  trickles  down  the  side  of  the  cathode  and  is  drawn  off 
from  beneath  the  oil.  The  constant  removal  of  the  hydroxide 
enables  the  electrolysis  to  be  carried  p/actically  to  completion 
and  the  yield  approaches  the  theoretical.  A  high  current  density 
is  employed. 

The  CASTNER  process  is  the  most  extensively  used  of  the 
amalgamation  type.  Its  cell  is  divided  into  three  compartments, 
the  two  outside  ones  containing  brine  and  the  carbon  anodes, 
while  the  middle  one  contains  the  caustic  liquor.  A  layer  of 
mercury  covers  the  bottom  of  the  whole  cell.  In  the  brine  com- 
partments the  mercury  acts  as  cathode,  taking  up  the  sodium  to 
form  amalgam.  The  amalgamated  mercury  is  transferred  to 
the  middle  compartment  where  it  is  decomposed  by  water  to 
form  sodium  hydroxide.  This  caustic  solution  is  drawn  off  and 
fresh  water  introduced  at  a  regulated  rate.  On  evaporation  a 
very  pure  sodium  hydroxide  results.  The  efficiency  of  the  proc- 
ess is  enhanced  by  conserving  the  electrical  energy  liberated  in 
the  decomposition  of  the  amalgam. 

In  one  successful  bell  process  the  anode  consists  of  some 
carbon  supported  in  a  bell  which  is  suspended  in  the  brine 
and  has  an  exit  tube  at  the  top  for  piping  off  the  chlorine. 
The  cathode  is  a  cylindrical  piece  of  sheet  iron  encircling  the 
bell. 

For  critical  discussions  of  the  relative  merits  of  these  and 


225.]  SALTS  OF  SODIUM.  315 

competitive  processes  the   reader  must   refer  to  the  technical 
journals  or  the  most  recent  works  on  industrial  chemistry. 

The  hydroxide  is  obtained  by  evaporation,  whereupon  it  is 
generally  cast  into  sticks  for  the  market.  It  is  radiate-crys- 
talline and  very  hygroscopic.  It  dissolves  in  water  with  the 
evolution  of  considerable  heat.  Sodium  hydroxide  is  a  very 
strong  base;  it  is  used  in  the  arts  for  numerous  purposes,  among 
others  the  manufacture  of  soap. 


SALTS  OF   SODIUM. 

225.  The  sodium  salts  are  of  great  industrial  importance; 
many  of  them  are  prepared  in  enormous  quantities.  The  starting- 
point  for  their  manufacture  is  usually  common  salt. 

Sodium  chloride,  NaCl,  common  salt,  is  found  in  large  masses 
as  rock-salt,  e.g.  at  Stassfurt  and  Reichenhall  and  in  Galicia,  where 
it  ,is  dug  out  by  miners.  Farther,  large,  amounts  are  obtained 
from  sea  water  and  the  water  of  salt  wells.  Three  methods  are 
employed  to  remove  salt.  In  sufficiently  warm  countries  (e.g. 
Mediterranean  coast,  central  New  York  State)  the  brine  is  led 
into  flat  basins  of  very  large  surface  area  (" salterns,"  or  "salt 
covers  ").  In  these  the  water  is  removed  by  solar  evaporation 
and  the  salt  crystallizes  out.  Any  gypsum  that  may  be  present 
separates  out  first,  whereupon  the  brine  passes  to  further  basins 
and  yields  pure  salt.  Later  on,  the  other  salts  separate  out; 
these  are  sometimes  worked  up  commercially.  In  countries 
with  a  cold  climate  (e.g.  on  the  shores  of  the  White  Sea)  the 
water  is  allowed  to  freeze  in  flat  basins.  The  ice  that  forms 
is  free  from  salt  so  that  the  remaining  liquid  is  more  con- 
centrated. 

In  countries  of  the  temperate  zone  the  sea  water  is  con- 
centrated by  letting  it  evaporate  spontaneously  from  a 
greatly  enlarged  surface.  This  is  done  by  the  "graduation" 
process  (Fig.  50).  Bundles  of  fagots  are  piled  up  together 
in  a  "rick,"  above  which  a  trough  with  small  outlet-holes 
runs  from  end  to  end.  The  brine  is  pumped  up  into  the 
trough  and  trickles  down  from  along  the  entire  length  of 
the  latter  upon  the  brush;  in  this  way  the  surface  of  the 
salt  solution  is  greatly  enlarged  and  the  evaporation  is  made 


316 


INORGANIC  CHEMISTRY. 


[§225 


much  more  rapid.     A  very  concentrated  brine  flows  out  at  the 
bottom. 

The  salt  is  obtained  from  this  concentrated  solution  by  boiling 
("salt-boiling").  Common  salt  is  almost  equally  soluble  in  hot 
and  cold  water,  hence  it  does  not  crystallize  out  on  cooling  but  falls 
out  at  the  same  rate  as  the  saturated  brine  evaporates,  even  while 


FIG.  50. — GRADUATION  PROCESS. 

boiling  hot.  The  salt  obtained  from  the  first  crystallization  is  of 
course  impure,  containing  small  amounts  of  magnesium  salts, 
which  render  it  hygroscopic.  In  order  to  purify  it,  it  is  redissolved 
in  water  and  again  precipitated  by  evaporation. 

Chemically  pure  sodium  chloride  is  obtained  by  passing  hydro- 
chloric acid  gas  into  a  saturated  solution  of  the  salt  or  treating  the 
solution  with  the  concentrated  acid.  The  sodium  chloride  is 
deposited  because  it  is  less  soluble  in  hydrochloric  acid  than  in 
water  (§  205). 

Common  salt  crystallizes  in  cubes;  when  the  solution  evaporates 


§  225.]  SALTS  OF  SODIUM.  317 

slowly  the  well-known  hollow  four-sided  pyramids,  or  hopper- 
crystals  (Fig.  51),  are  formed.  Sp.  g.=2.16.  M.-pt.=776°. 
100  parts  of  water  dissolve  36  parts  NaCl  at  0°. 
39  parts  at  100° ;  a  saturated  solution  contains 
about  26%  of  salt.  The  crystals  frequently 
enclose  some  of  the  mother-liquor;  for  this 
reason  they  decrepitate  on  heating.  On  cool- 
ing below  —10°  a  saturated  solution  depos- 
its crystals  of  the  composition  NaCl  +  H2O, 
which  lose  their  water  at  0°.  Chemically 
pure  sodium  chloride  is  not  hygroscopic.  It 
i8  insoluble  in  absolute  alcohol. 

Sodium  bromide,  NaBr,  and  sodium  iodide,  Nal,  are  more  soluble 
in  water  than  the  chloride.  From  hot  solutions  they  crystallize  in 
anhydrous  cubes,  below  30°  in  monoclinic  crystals  with  2  mols.  H20. 
Sodium  bromide  is  difficultly  soluble,  sodium  iodide  easily  soluble,  in 
alcohol. 

Sodium  thiosulphate,  Na2S203-5H20;  spoken  of  sometimes  as 
sodium  hyposulphite  or  "  hypo/'  is  employed  in  photography 
(§  247) ,  in  titrating  iodine  (§  93),  and  as  an  antichlor  (§  82). 
It  melts  in  its  water  of  crystallation  at  45°  and  forms  super- 
saturated solutions  very  easily. 

Sodium  sulphate,  Na2SO4-10H20  (sal  mirabile  Glauberi,  Glau- 
ber's salt — so  called  after  the  discoverer),  can  be  obtained  in  various 
ways:  (1)  by  heating  common  salt  with  concentrated  sulphuric 
acid;  (2)  by  conducting  a  mixture  of  air,  sulphur  dioxide  and 
steam  over  hot  sodium  chloride  (HARGREAVE'S  method): 

2NaCl + S02  +  O  +  H2O  =  Na2SO4 + 2HC1  ; 

(3)    by   the   double   decomposition   of   magnesium   sulphate   and 
sodium  chloride  at  a  low  temperature  (winter- temperature) : 

MgSO.i + 2NaCl = MgCl2  +  Na2SO4. 

This  last  process  is  carried  out  at  Stassfurt,  where  large  masses  of 
magnesium  sulphate  occur. 

At  ordinary  temperatures  Glauber's  salt  crystallizes  with  ten 
molecules  of  water ;  above  33°  it  goes  over  into  a  mixture  of  water 


318  INORGANIC  CHEMISTRY.  [§§  225- 


and  anhydrous  salt;  below  33°  the  hydrous  salt, 

is  again  formed.      The  system  Na2SO4-10H2O^Na2SO4-}-10H2O 

thus  has  a  transition  point  at  33°.     This  is  confirmed  by  the  fact 

that  the  solubility  of  Glauber's  salt  in  water  suddenly  changes  at 

33°  from  a  rapid  increase  with  rising  temperature  to  a  slow  de- 

crease. 

This  salt,  like  the  preceding  one,  forms  supersaturated  solutions 
easily.  The  solution  saturated  at  33°  can  be  cooled  down  to  room- 
temperature  without  any  deposition  if  care  is  taken  to  exclude 
even  the  tiniest  crystal  of  the  salt. 

When  a  crystal  of  Glauber's  salt  is  exposed  to  the  air  it  effloresces, 
i.e.,  it  gives  off  water  vapor  and  becomes  opaque,  like  chalk.  This 
is  evidently  due  to  the  fact  that  the  vapor  tension  of  its  water  of 
crystallization  is  greater  throughout  than  that  of  the  water  vapor 
in  the  air.  Inversely,  we  say  that  a  salt  deliquesces  when  it  takes 
up  water  vapor  from  the  air,  as  a  result  of  the  vapor  tension  of 
its  saturated  solution  being  less  than  the  mean  tension  of  the  water 
vapor  in  the  atmosphere.  It  is  found  that  a  perfectly  sound  crystal 
of  Na2SO-4  •  10H20  does  not  effloresce,  but  that  when  efflorescence 
has  begun  at  any  point  it  spreads  over  the  crystal.  The  phase  rule 
gives  a  satisfactory  explanation  of  this  phenomenon.  We  have 
in  the  Glauber's  salt  two  substances,  Na2S04  and  H20;  in  the  case 
of  a  perfectly  bright  crystal  exposed  to  the  air  we  have  only  two 
phases,  Na2S04  •  10H2O  and  H2O  (moisture  of  the  air).  According 
to  the  equation  F+P=S+2  (§71)  we  still  have  in  this  system 
two  degrees  of  freedom,  i.e.,  the  pressure  of  the  water  vapor  and 
the  temperature  can  both  be  selected  arbitrarily  (within  certain 
limits).  If,  however,  some  dehydrated  salt  is  present,  the  number 
of  phases  is  three;  then  there  is  only  one  degree  of  freedom,  or; 
in  other  words,  every  temperature  has  only  one  corresponding 
pressure  and  inversely  every  pressure  only  one  corresponding 
temperature.  Accordingly  it  is  only  permissible  to  speak  of  the 
well-defined  vapor  tension  of  a  salt  with  water  of  crystallization 
when  a  second  solid  phase  is  tacitly  admitted  to  be  present;  for 
then  only  is  the  number  of  degrees  of  freedom  reduced  to  one. 

Glauber's  salt  is  used  in  medicine;  in  the  arts  it  is  employed 
chiefly  in  the  manufacture  of  soda. 

Sodium  nitrate,  NaNO3,  called  Chili  saltpetre  because  of  its 
extensive  occurrence  in  Chili,  crystallizes  in  rhombohedrons  and 


223.]  SALTS  OF   SODIUM.  319 

melts  at  318°.  It  is  somewhat  hygroscopic.  Large  quantities  of 
it  are  used  in  fertilizing  and  also  in  the  preparation  of  nitric  acid 
and  potash  saltpetre. 

Sodium  nitrite,  NaNO2,  is  obtained  by  heating  the  nitrate,  the 
addition  of  a  reducing-agent  such  as  lead  or  iron  aiding  the  process. 
It  is  very  soluble  in  water  and  is  consumed  in  large  quantities  in 
the  aniline-dye  industry. 

Sodium  phosphates.  (See  §  146.)  Trisodium  phosphate, 
Na3P04  •  12H20,  is  split  up  in  aqueous  solution  chiefly  into 
sodium  hydroxide  and  the  secondary  salt,  for  the  solution  reacts 
strongly  alkaline  and  absorbs  carbonic  acid  from  the  air.  The 
ordinary  "sod'.um  phosphate"  is  the  disodium  phosphate, 
Na2HPO4  •  12H2O.  It  separates  from  its  aqueous  solution  at 
ordinary  temperatures  in  large  crystals  which  soon  effloresce.  100 
parts  H2O  dissolve  9.3  parts  of  the  salt  at  20°  and  24.1  parts  at 
30°.  The  solution  is  feebly  alkaline.  By  leading  in  carbonic  acid 
gas  a  liquid  of  amphoteric  reaction  (cf.  §  224)  is  obtained,  which 
turns  blue  litmus  red  as  well  as  red  litmus  blue.  Monosodium 
phosphate,  NaH^PCVRsO,  reacts  acid.  It  is  converted 
by  heat  into  the  metaphosphate. 

226.  Sodium  carbonate,  Na2CO3-10H2O  (soda,  sal-soda),  is, 
next  to  the  chloride,  the  most  important  sodium  salt  and  it  is 
manufactured  on  an  enormous  scale.  It  occurs  in  nature  in  Egypt, 
in  America  (Wyoming),  on  the  Caspian  Sea,  and  elsewhere;  the 
ashes  of  many  marine  plants  contain  it. 

The  manufacture  of  soda  is  carried  on  mainly  by  two  different 
methods : 

(1)  The  LE  BLANC  soda  process. 

This  process  consists  of  three  parts.  In  the  first  place  common 
salt  is  warmed  with  strong  sulphuric  acid  (chamber  acid) ; 
hydrochloric  acid  and  sodium  sulphate  "  salt  cake  "  are  formed. 
In  the  second  part  of  the  process  sodium  sulphate  is  heated  with 
coal  and  limestone.  The  third  section  of  the  process  consists 
in  lixiviating  the  mass  last  obtained  ("  black-ash  ")  with  water, 
whereby  sodium  carbonate  is  dissolved  out.  The  latter  is  then 
obtained  from  this  solution  by  crystallization. 

After  the  black  ash  has  been  leached  out  as  far  as  practicable, 
it  is  cast  aside  as  "  tank  waste."  The  most  valuable  constituent 
of  the  latter  is  calcium  sulphide ;  it  is  conserved  by  treating  the 


320  INORGANIC  CHEMISTRY.  (§  226. 

waste  with  water  and  carbon  dioxide  to  form  hydrogen  sulphide, 
which  is  then  oxidized  to  sulphur  and  the  latter  to  sulphuric 
acid.  The  process  in  its  entirety  is  thus  represented  by  the 
following  equations: 

.  2NaCl  +  H2SO4=Na2SO4+2HCl; 

Na2SO4+  2C  =  2C02  +  Na2S  ; 
Na2S  +  CaCO3  =  CaS  +  Na2CO3  ; 
G02  +  H2O  +  CaS  =  CaC03  +  H2S  ; 
=  H2SO4; 


or,  summed  up  : 

2NaCl  +  CO2  +  H2O  =  2HC1  +  Na2CO3. 

The  process  is  noted  for  its  high  efficiency,  since  all  the  by-products 
are  worked  up.  Nevertheless,  this  process,  which  for  a  long  period  of 
years  practically  controlled  the  industrial  market,  is  now  almost  wholly 
superseded  by  the  other  one  and  a  few  years  more  will  probably  see  its 
entire  abandonment. 

(2)  The    ammonia-soda    process    ofSoLVAY. 

This  process,  which  originally  presented  numerous  technical 
difficulties,  is  now  so  perfected  that  about  ninety-five  per  cent 
of  the  total  soda  production  is  by  the  SOLVAY  process.  The 
chemistry  of  the  SOLVAY  process  is  very  simple.  Ammonia 
and  carbon  dioxide  are  led  alternately  into  a  cold  concentrated 
salt  solution  under  pressure.  The  following  reaction  then  takes 
place  : 

NaCl  +  (NH4)  HCO3  =  NaHCO3  +  NH4C1. 

The  acid  sodium  carbonate  ("  bicarbonate  ")  so  formed  sepa- 
rates out,  inasmuch  as  it  is  very  difficultly  soluble  in  the  cold 
concentrated  ammonium  chloride  solution.  It  is  broken  up,  on 
heating,  into  soda  and  carbon  dioxide,  the  latter  of  which  is 
carried  back  to  be  used  over.  The  ammonium  chloride  solution 
is  distilled  with  lime,  whereby  ammonia  is  recovered. 

The  process  as  a  whole  may  be  represented  by  the  following 
equations  : 


226.]  SODIUM  CARBONATE.  321 

2NaCl  +  2NH3  +  2CO2  +  2H2O  =  2NH4C1  +  2NaHCO3  ; 
2NaHC03  =  H2O  +  CO2  +  Na2CO3  ; 
2NH4C1  +  CaO  =  2NHg  +  H2O  +  CaCl2  ; 
CaC03  = 


or,  summed  up  : 

2NaCl  +  CaC03  =  Na2C03  +  CaCl2. 

In.  the  SOLVAY  process  there  is  formed  together  with  the  soda  an 
equivalent  amount  of  calcium  chloride,  for  which  there  is  only  a  limited 
market  (chiefly  for  refrigeration;  cf.  §  258),  so  that  one  valuable  con- 
stituent of  salt,  the  chlorine,  is  largely  lost.  Attempts  to  substitute 
magnesia  for  the  lime  so  as  to  be  able  to  utilize  the  resulting  magnesium 
chloride  for  the  recovery  of  hydrochloric  acid  or  chlorine  have  not  been 
commercially  successful. 

Some  soda  is  also  manufactured  by  carbonating  the  electrolyt- 
ically  prepared  sodium  hydroxide. 

Sodium  carbonate  crystallizes  at  ordinary  temperatures  with 
ten  molecules  of  water  of  crystallization  in  large  transparent 
monoclinic  crystals,  which  soon  turn,  white  and  dull  from  loss  of 
water  (efflorescence).  They  melt  at  60°  in  their  own  water;  on 
continued  warming  the  hydrate  Na2CO3  +  2H2O  is  deposited;  the 
latter  loses  one  molecule  of  water  in  dry  air  and  the  remainder  at 
100°.  At  30°-50°  rhombic  prisms  of  the  composition  Na2CO3 
+  7H20  crystallize  out  of  the  aqueous  solution.  100  parts  H2O 
dissolve  6.97  parts  of  the  anhydrous  salt  at  0°,  51.67  parts  at  38°. 
The  aqueous  solution  reacts  strongly  alkaline  (§  184)  because  ' 
hydrolysis.  As  was  set  forth  in  §  146,  this  phenomenon  m, 
always  occur  when  a  salt  of  a  weak  base  and  a  strong  acid  or  u. 
salt  of  a  weak  acid  and  a  strong  base  is  formed.  In  case  both 
acid  and  base  are  weak  the  hydrolysis  will  be  all  the  greater. 
Which  reaction  the  solution  of  such  a  salt  will  give  depends  on 
the  relative  strengths  of  the  acid  and  base. 

Soda  is  used  commercially  on  a  large  scale,  particularly  in 
the  soap  and  glass  industries.  It  is  the  "  washing-soda  "  of  the 
household. 


322  INORGANIC  CHEMISTRY.  [§§226- 

Acid  sodium  carbonate,  NaHC03  (bicarbonate  of  soda),  is 
obtained  as  a  primary  product  in  the  SOLVAY  process.  It  dis- 
solves in  10-11  parts  of  water  at  room  temperature  and  reacts 
alkaline.  On  being  gently  warmed  it  breaks  up  into  carbon 
dioxide,  water  and  soda;  this  decomposition  occurs  even  on  warm- 
ing the  aqueous  solution,  and  when  a  current  of  air  is  passed 
through  the  concentrated  solution  at  ordinary  temperatures  car- 
bon dioxide  escapes.  It  is  used  extensively  in  baking-powders 
and  is  the  saleratus  of  commerce. 

Sodium  silicate  (sodium  water-glass)  is  prepared,  among  other 
ways,  by  fusing  sand  with  Glauber's  salt  and  charcoal.  This 
yields  a  vitreous  mass,  which  is  dissolved  by  boiling  water.  The 
concentrated  solution  has  the  consistency  of  glue.  It  finds  use  as 
a  fixative  in  calico  printing,  as  well  as  for  impregnating  inflam- 
mable textiles  like  theater  decorations,  etc.;  it  is  also  used  for 
"  filling  "  soaps. 

The  sulphides  of  sodium  correspond  to  those  of  potassium  and 
are  prepared  in  the  same  way  (see  §  231). 

Sodium  borate:   cf.  Borax  (§283). 


POTASSIUM. 

227.  Compounds  of  potassium  occur  in  nature  very  extensively 
but  not  in  such  large  quantities  as  those  of  sodium.  Potassium 
exists  principally  in  the  silicates,  especially  feldspar  and  mica. 
Upon  the  decay  of  these  minerals  it  is  carried  into  the  soil  and 
thence  into  the  plants,  to  which  potassium  compounds  are  indis- 
pensable. Potassium  salts  are  also  found  in  sea- water.  The 
largest  source  of  them,  however,  is  the  Stassfurt  "  Abraum 
salts  "  (§  44),  mainly  double-salts  of  potassium  and  magnesium, 
such  as  carnallite,  MgCl2  •  KC1  •  6H20,  kainite,  MgS04  •  Kd  •  3H2O, 
etc.  The  large  amounts  of  potassium  in  the  feldspars  makes  its 
recovery  from  them  a  very  enticing  problem. 

The  metal  was  first  obtained  by  DAVY  by  the  electrolysis  of 
molten  caustic  potash.  One  of  the  commercial  methods  is  to 
ignite  a  mixture  of  carbonate  of  potash  and  powdered  charcoal 
(preferably  charred  acid  potassium  tartrate).  The  extraction  of 
the  metal  is  thus  analogous  to  that  of  sodium;  in  the  preparation 
of  potassium,  however,  potassium  carbonyl,  Co(OK)6,  may  be 


228.  ]  OXYGEN  COMPOUNDS  OF  POTASSIUM.  323 

formed  under  certain  circumstances,  a  substance  which  acquires 
explosive  properties  on  exposure  to  the  air. 

Potassium  has  a  silvery-white  metallic  lustre  and  is  almost  as 
soft  as  wax  at  ordinary  temperatures.  Sp.  g.  =0.875  at  13°.  It 
melts  at  62.5°  and  boils  at  about  720°,  forming  a  green  vapor.  The 
mirror-like  surface  of  the  metal  immediately  becomes  dull  in  the 
air;  when  heated  in  the  air  it  burns  with  an  intense  violet  light. 
Water  is  decomposed  by  it  with  great  vigor,  the  heat  evolved 
being  sufficient  to  ignite  the  escaping  hydrogen  and  drive  the 
piece  of  potassium  around  on  the  water. 

Oxygen  Compounds  of  Potassium. 

228.  Potassium  oxide,  K2O,  is  formed  by  oxidizing  potas- 
sium by  the  method  described  in  §  224.  It  is  a  white  substance, 
which  unites  with  water  to  form  the  hydroxide  with  the  evolution 
of  much  heat. 

Potassium  peroxide,  KC>2,  is  produced  together  with  the  mon- 
oxide on  burning  potassium  in  the  air.  It  is  dark  yellow.  In  con- 
tact with  water  it  yields  potassium  hydroxide,  hydrogen  peroxide 
and  free  oxygen. 

Potassium  hydroxide,  KOH,  results  from  the  action  of  potas- 
sium on  water  and  is  generally  prepared  in  the  same  manner  as 
sodium  hydroxide,  viz.,  by  treating  potassium  carbonate  solution 
with  milk  of  lime,  Ca(OH)2.  It  can  also  be  obtained  by  heating 
saltpetre  with  powdered  copper  (forming  copper  oxide  and  potas- 
sium oxide),  and  adding  water;  the  copper  oxide  can  be  removed 
by  filtration.  The  hydroxide  usually  comes  on  the  market  in 
sticks. 

The  commercial  product  ("  caustic  potash  ")  is  obtained  chiefly 
by  the  first  method  and  usually  contains  a  little  sulphate,  chloride, 
etc.,  besides  the  carbonate  which  is  gradually  formed  by  the  action 
of  atmospheric  carbon  dioxide.  It  can  be  purified  by  treating  with 
strong  alcohol,  which  dissolves  only  the  hydroxide;  after  filtering, 
the  alcoholic  solution  is  evaporated  in  a  silver  dish.  Caustic  soda 
is  also  purified  in  this  way. 

Potassium  hydroxide  is  one  of  the  strongest  bases.  In  the  solid 
state  it  greedily  absorbs  water  and  carbon  dioxide  from  the  air  and 
finally  deliquesces  to  a  concentrated  solution  of  potassium  carbonate, 
while  sodium  hydroxide  under  these  conditions  turns  to  a  solid 


324  INORGANIC  CHEMISTRY.  [§§228- 

white  mass  of  soda.  For  this  reason  caustic  potash  is  a  much  more 
suitable  absorptive  agent  for  carbon  dioxide  in  analyses  than  caustic 
soda,  for  the  use  of  the  latter  might  easily  cause  a  stopping  up  of 
the  apparatus. 

Caustic  potash  is  used  especially  in  the  manufacture  of  soft 
soaps. 

Potassium  Salts. 

229.  Potassium  chloride,  KC1,  occurs  at  Stassfurt  in  the  min- 
eral sylvite.  It  crystallizes  in  cubes  and  melts  at  730°.  It  is  easily 
volatilized  at  elevated  temperatures.  100  parts  H2O  dissolve  25.5 
parts  KC1  at  0°,  57  parts  at  100°.  Like  its  sodium  analogue, 
potassium  chloride  is  precipitated  from  its  saturated  solution  by 
hydrochloric  acid.  It  unites  with  many  salts  to  form  double  salts. 

Potassium  bromide,  KBr,  is  important  therapeutically.  It  is 
prepared  by  mixing  bromine  with  a  potassium  hydroxide  solution, 
the  bromide  and  bromate  being  formed;  the  bromate  is  reduced 
by  heating  the  salty  product  with  powdered  charcoal.  Potassium 
bromide  crystallizes  in  cubes  and  dissolves  readily  in  water. 

Potassium  iodide,  KI,  also  of  medicinal  value,  can  be  prepared 
like  the  bromide  and  also  in  the  following  manner:  Iodine  and 
iron  filings  are  mixed  together  under  water,  whereupon  a  solu- 
tion of  the  compound  Feslg  is  formed;  on  treating  this  with  a 
potash  solution  the  oxide  Fe304  is  precipitated,  carbon  dioxide 
escapes  and  potassium  iodide  is  left  in  solution;  the  salt  is  then 
obtained  by  filtration  and  evaporation.  It  crystallizes  in  cubes 
and  is  very  soluble  in  water:  1  part  H20  dissolves  1.278  parts 
KI  at  0°.  On  exposure  to  light  or  the  air  the  crystals  gradually 
turn  yellow  because  of  the  separation  of  iodine. 

It  was  remarked  in  §  46  that  iodine,  though  only  slightly 
soluble  in  water,  dissolves  to  a  much  greater  extent  when  the 
water  contains  potassium  iodide.  This  is  due  to  the  formation 
of  JV  ions  in  the  latter  case.  That  the  iodine  has  entered  into 
combination  may  be  concluded  in  the  first  place  from  the  fact 
that  the  addition  of  iodine  to  an  aqueous  solution  of  potassium 
iodide  does  not  cause  a  further  depression  of  the  freezing-point; 
the  number  of  molecules  is  thus  unchanged,  or,  in  other  words, 
iodine  has  combined  with  potassium  iodide:  in  the  second  place, 
from  the  fact  that  carbon  disulphide  takes  up  nearly  all  tHe  iodine 
from  an  aqueous  solution  of  the  latter  when  it  is  shaken  with  the 
solution,  but  only  a  small  proportion  when  the  same  operation  is 


229.  POTASSIUM  SALTS.  325 

performed  with  a  solution  of  iodine  in  a  dilute  aqueous  solution  of 
potassium  iodide.  The  distribution  ratio  for  iodine  between  water 
and  carbon  disulphide  is  1:410.  If,  therefore,  we  divide  the  con- 
centration of  the  iodine  in  carbon  disulphide  by  410  we  obtain  the 
concentration  of  the  free  iodine  in  the  potassium  iodide  solution. 
Subtracting  this  from  the  total  concentration  of  the  iodine  in 
this  solution,  we  have  the  amount  of  combined  iodine.  It  is  found 
that  1KI,  or,  rather,!!',  has  taken  up  1I2.  Is'  ions  are  thus  formed 
in  .the  solution. 

Nevertheless,  a  solution  of  iodine  and  potassium  iodide  in 
water  behaves  in  many  cases  as  if  all  the  iodine  were  present  in 
the  free  state,  e.g.  when  it  is  titrated  with  sodium  thiosulphate. 
This  must  be  explained  by  the  supposition  that  in  the  liquid  we 
have  the  equilibrium: 


If  the  free  iodine  is  removed,  the  equilibrium  is  disturbed;  a 
new  portion  of  1%  must  therefore  split  up,  and  so  on  till  it  is  entirely 
consumed. 

Potassium  fluoride,  KF,  possesses  a  peculiar  property,  which  is 
lacking  with  the  other  halogen  compounds  of  potassium:  it  com- 
bines with  hydrofluoric  acid  eagerly,  forming  the  double  halide 
KF-HF. 

Potassium  cyanide,  KCN  (often  also  written  KCy),  is  manufac- 
tured on  a  large  scale  by  fusing  yellow  prussiate  of  potash  with 
potash  : 

K4Fe(CN)6  +  K2C03  =  5CNK  +  KCNO  +  C02  +  Fe. 

The  cyanate  of  potassium  KCNO  is  reduced  by  the  iron  to  potassium 
cyanide  also.  It  is  very  soluble  in  water,  forming  a  strongly 
alkaline  solution.  On  account  of  its  great  tendency  to  form  double- 
salts,  it  is  employed  in  electro-metallurgy.  It  is  also  used  in 
extracting  gold  from  its  ores  (§  248). 

Potassium  chlorate,  KClOs,  can  be  obtained  by  passing  chlo- 
rine into  a  h  o  t  solution  of  caustic  potash  (§  56).  It  is  now  pre- 
pared almost  exclusively  by  the  electrolysis  of  a  hot  solution  of 
sodium  chloride.  If  in  the  GRIESHEIM  process  (§  224)  for  the 
manufacture  of  caustic  alkali  the  electrolysis  is  continued  after 
a  certain  amount  of  hydroxide  has  formed,  the  oxygen  liberated 


326  INORGANIC  CHEMISTRY.  [§§  229- 

at  the  anode  oxidizes  the  sodium  chloride  to  chlorate,  NaClO3. 
The  latter  is  converted  into  the  potassium  chlorate  by  treatment 
with  potassium  chloride.  The  advantage  of  this  method  is  that 
sodium  chlorate  is  much  more  soluble  and  does  not  retard  the 
electrolytic  process  by  separating  from  the  solution,  as  does 
potassium  chlorate. 

Potassium  chlorate  is  a  well-crystallized  salt,  which  is  used  for 
the  preparation  of  oxygen  (§  9) ;  furthermore,  it  is  used  in  the 
manufacture  of  matches  and  fireworks,  and  also  medicinally  as  a 
remedy  for  sore  throat.  On  being  heated  it  gives  up  oxygen, 
part  of  the  salt  being  at  the  same  time  converted  into  potassium 
perchlorate,  KC104. 

The  last-named  salt  is  difficultly  soluble  in  water.  It  is  sometimes 
found  in  crude  Chili  saltpetre,  rendering  the  latter  unfit  for  use  in 
fertilizing  various  cultivated  plants. 

Potassium  sulphate,  K2S04,  is  obtained  by  the  action  of  sul- 
phuric acid  on  potassium  chloride.  It  crystallizes  in  beautiful, 
lustrous  rhombic  prisms  and  dissolves  with  some  difficulty  in  cold 
water  (1  part  in  10  parts  H20  at  room  temperature).  It  is  used 
principally  for  the  preparation  of  potash  according  to  the  LE 
BLANC  method.  Acid  potassium  sulphate,  KHS04,  is  very  soluble 
in  water;  it  melts  at  200°,  losing  water  and  going  over  into  potas- 
sium pyrosulphate,  K2S2O7.  The  latter  breaks  up  into  potassium 
sulphate  and  sulphur  trioxide  on  heating. 

Potassium  nitrate,  KNOs,  is  widely  distributed  in  nature, — 
although  usually  found  only  in  small  amounts, — for  it  is  formed 
wherever  nitrogeneous  organic  bodies  decay  in  contact  with  potas- 
sium compounds.  This  is  the  basis  of  an  artificial  method  of 
preparing  saltpetre,  which  method  was  formerly  much  used. 

Another  process  of  manufacture  depends  on  the  double  decom- 
position of  Chili  saltpetre  with  potassium  chloride,  which  is  obtained 
in  large  quantities  at  Stassfurt: 

KC1 + NaNO3  =  KN03  +  NaCl. 

For  this  purpose  hot-saturated  solutions  of  the  two  salts  are  brought 
together.  As  sodium  chloride  is  much  less  soluble  than  saltpetre 


231.]  POTASSIUM   SALTS.  .       327 

at  the  temperature  of  boiling  water,  it  is  the  first  to  crystallize 
out  on  evaporation,  but  when  the  solution  is  cooled  the  saltpetre 
comes  out  first,  for  it  is  much  less  soluble  than  sodium  chloride 
in  cold  water. 

Potash  saltpetre  crystallizes  in  anhydrous  prisms,  either  rhom- 
bohedral  or  rhombic  according  to  the  temperature.  In  the  neigh- 
borhood of  the  melting-point  the  former  is  the  stable  variety 
at  ordinary  temperatures  the  latter.  The  location  of  the  transi- 
tion point  of  the  two  forms  has  not  yet  been  determined.  100 
parts  H20  dissolve  13.3  parts  KN03  at  0°,  247  at  100°.  It  melts 
at  338°;  farther  heating  breaks  it  up  into  potassium  nitrite  and 
oxygen.  It  has  a  cooling  taste. 

230.  Potash  saltpetre  is  consumed  in  large  quantities  in  the 
manufacture  of  gunpowder.     This  is  a  mixture  of  sulphur,  charcoal 
and  potash  saltpetre,  the  proportions  varying  in  different  countries, 
but  being  in  most  cases  75%  KN03,  10%  S  and  15%  charcoal. 

231.  Potassium  phosphates. — The    three    potassium    salts    of 
phosphoric  acid  are  known.     They  are  very  soluble  in  water. 

Potassium  carbonate,  K2CO3,  potash. — This  salt  was  formerly 
obtained  solely  from  wood-ashes,  these  being  soaked  in  water  and 
the  strained  liquor  evaporated.  At  present  it  is  manufactured 
extensively  from  potassium  chloride  after  the  LE  BLANC  process. 
Another  source  of  potash  is  the  molasses  of  the  beet-sugar  fac- 
tories, that  contains  the  potassium  salts  in  which  the  sugar-beet  is 
rich. 

At  the  Neustassfurt  salt  mine  it  is  made  from  potassium  chloride 
by  a  patented  process  as  follows:  Magnesium  carbonate,  MgC03-3H20, 
is  suspended  in  a  solution  of  potassium  chloride,  and  carbon  dioxide 
is  led  in,  whereupon  the  following  reaction  takes  place: 

3MgC03.3H2O  +2KC1 +C02  =MgCl2  +  2MgC03  •  KHC03  •  H20. 

The  potassium  magnesium  carbonate  separates  out  and  is  broken  up 
by  heating  to  a  temperature  not  exceeding  80°  into  magnesium  car- 
bonate and  potash.  The  former  salt  is  again  obtained  with  three  mole- 
cules of  water  of  crystallization,  which  form  is  the  only  one  suited  for 
the  above  reaction. 

Potassium  carbonate  is  a  white  powder,  which  deliquesces 
in  the  air  and  is  very  soluble  in  water  (1.12  parts  K2COs  in  1  part 


328  INORGANIC  CHEMISTRY.  [§§  231- 

H2O  at  20°) ;  the  solution  has  a  strong  alkaline  reaction.  The  salt 
melts  at  838°.  It  is  used  in  the  preparation  of  soft  soaps  and  hard 
glass  (potash-glass). 

Potassium  silicate,  potassium  water-glass,  is  formed  when  sand 
is  fused  with  potash.  Different  salts  of  this  sort  are  described. 
They  dissolve  in  water,  forming  a  thick,  mucilaginous  mass  which 
on  drying  turns  to  a  vitreous,  and  finally  opaque,  product. 
Potassium  water-glass  is  used  for  the  same  purposes  as  sodium 
water-glass. 

Sulphides  of  Potassium. 

Potassium  monosulphide,  K2S,  is  prepared  by  reducing 
potassium  sulphate  with  charcoal.  It  dissolves  in  water  very 
readily  and  crystallizes  out  with  five  molecules  of  water. 
It  absorbs  oxygen  from  the  air;  going  over  into  the  thio- 
sulphate  and  hydroxide: 

2K2S  +  H20  +  202  =  K2S203  +  2KOH. 

Acids  react  with  it,  liberating  hydrogen  sulphide. 

Potassium  hydrosulphide,  KSH,  is  obtained  by  saturating  a 
caustic  potash  solution  with  hydrogen  sulphide: 

KOH+H2S=KSH  +  H20. 

It  is  very  soluble  in  water,  the  solution  reacting  alkaline;  on 
evaporation  in  vacuo  the  solution  deposits  crystals  of  the  com- 
position 2KSH  +  H2O.  With  potassium  hydroxide  it  forms  the 
monosulphide : 

KSH+KOH=K2S  +  H20. 

Potassium  polysulphides. — When  a  solution  of  potassium  mono- 
sulphide  is  boiled  with  sulphur,  we  obtain  the  compounds  K2Ss, 
K2$4,  K2S5.  A  mixture  of  these  substances  is  also  obtained  by 
fusing  potash  with  sulphur;  besides  these  it  contains  the  sulphate 
and  the  thiosulphate  and  is  called  hepar  sulphuris  ("  liver  of  sul- 
phur ")  because  of  its  liver-brown  color.  These  polysulphides  are 


232.]  RUBIDIUM  AND  CESIUM.  329 

decomposed  by  acids  with  the  evolution  of  hydrogen  sulphide  and 
the  separation  of  sulphur: 

K2SX  +  2HC1  =  2KC1 + H2S  +  (x  - 1)8. 


Rubidium  and  Caesium. 

232.  These  elements  are  widely  distributed,  but  always  occur  in 
extremely  small  amounts.  The  silicate  lepidolite,  or  lithia  mica,  fre- 
quently contains  a  little  rubidium.  The  exceedingly  rare  mineral  pollux 
from  the  isle  of  Elba  is  a  silicate  of  aluminium  and  caesium,  and  contains 
about  30%  caesium  oxide.  In  general  these  elements  are  found  where- 
ever  potassium  salts  are  met  with:  in  mineral  springs,  in  the  Stassfurt 
salts  (carnallite  contains  rubidium),  etc.  They  were  discovered  by 
BUNSEN  and  KIRCHHOFF  in  1860  with  the  aid  of  spectrum  analysis 
(§  264)  and  obtained  their  names  from  the  most  important  lines  in 
their  spectra  (rubidus  =  d&Tk  red;  ccesius  =  sky-blue.)  The  spectrum 
lines  were  used  as  a  test  in  the  separation  of  these  elements  from  the 
others;  after  trying  a  possible  method  of  separation  the  two  sc,var:t3 
would  see  which  portion  showed  the  lines  of  these  elements  the  brightest ; 
this  portion  was  then  examined  further. 

In  order  to  separate  them  from  the  large  amount  of  potassium  salts 
with  which  they  generally  occur,  they  are  converted  into  chlorides 
and  the  solution  is  evaporated,  whereupon  the  dry  residue  is  extracted 
with  strong  alcohol.  Almost  all  the  sodium  chloride  and  potassium 
chloride  remains  behind,  while  the  chlorides  of  rubidium  and  caesium 
dissolve.  Platinum  chloride  is  then  added  to  precipitate  K2PtCl6, 
Rb2PtCle,  and  Cs2PtCl6;  the  solubility  of  these  double  salts  in  water 
is  quite  different  (at  10°  100  parts  H20  dissolve  0.9  parts  K2-salt,  0.154 
Rby-salt,  and  0.05  Cs2-salt),  so  that  they  can  be  very  well  separated 
by  fractional  extraction  with  boiling  water.  The  rubidium  iron  "alum" 
is  particularly  well  suited  for  the  purification  of  rubidium  salts,  and 
especially  for  their  separation  from  potassium  salts,  since  it  is  readily 
soluble  in  hot,  and  only  slightly  soluble  in  cold,  water,  and  moreover 
crystallizes  beautifully;  potassium  iron  alum,  on  the  other  hand,  is 
very  soluble  even  in  cold  water. 

The  metals  rubidium  and  caesium  are  best  obtained  by  heating  their 
hydroxides  with  calcium  filings  in  a  vacuum.  The  metals  then  distil 


330 


INORGA  NIC  CHE  MIS  TR  Y. 


[§§232- 


off.  Rubidium  has  a  silvery  lustre,  melts  at  38.5,  and  has  a  specific 
gravity  of  1.522  at  15°.  The  metal  oxidizes  very  rapidly  in  the  air  or 
in  oxygen,  forming  dark  brown  crystals  of  the  peroxide,  RbO2.  On 
being  heated  in  current  of  hydrogen  it  yields  the  hydroxide  and  free 
oxygen : 

2Rb02 + 2H2  =  2RbOH  +  H20  +  O. 

Rubidium  oxide,  when  prepared  in  the  same  way  as  Na20,  is  obtained 
as  transparent,  pale  yellow  crystals  which  turn  golden  yellow  on  heating 
but  lose  this  color  again  on  cooling. 

The  hydroxide  is  a  very  strong  base;  its  salts  show  much  similarity 
to  the  analogous  potassium  compounds;  they  are  in  several  instances 
less  soluble,  e.g.,  Rb-alum,  Rb-perchlorate  (§  60),  etc. 

Csesiumisa  silvery-white  metal;  sp.  g.  1.85;  m.-pt.  26.5°;  b.-pt. 
670°.  It  soon  takes  fire  on  exposure  to  the  air.  The  oxide,  Cs20, 
obtained  in  the  same  way  as  the  other  alkali  oxides,  is  crystallized  and 
is  orange-colored  at  room  temperature  but  almost  black  at  250°. 
The  salts  of  csesium  are  very  similar  to  those  of  rubidium ;  some  of  them 
are  even  less  soluble,  and  are  therefore  used  for  the  preparation  of 
pure  csesium  compounds.  This  is  particularly  true  of  the  platinum 
double-salt  already  mentioned  and  the  csesium  alum  and  the  acid 
tartrate. 

Rubidium  bromide  and  iodide,  and  even  more  so  the  corresponding 
compounds  of  caesium,  have  the  property  of  combining  with  two  atoms 
of  bromine  and  iodine,  forming  yellow  or  brown  crystalline  compounds, 
e.g.  CsI3;  these  metals  can  thus  be  trivalent. 

SUMMARY  OF  THE  GROUP  OF  ALKALI  METALS. 

233.  The  gradual  change  of  the  physical  properties  of  these 
metals  with  increasing  atomic  weight  is  made  plain  by  the  follow- 
ing table: 


Li 

Na 

K 

Rb 

Cs 

Atomic  weight  

7.00 

23.00 

39  10 

85.45 

132.81 

Specific  gravity 

0.59 

0.97 

0.865 

1.52 

1.85 

Melting-point            .  .  . 

180.0° 

97.6° 

62.5° 

38.5° 

26°-27° 

Boiling-point        

<1400° 

877° 

757° 

696° 

670° 

Atomic  volume  

11.8 

23.7 

45.3 

56.7 

71.9 

234.]  AMMONIUM  SALTS.  331 

The  specific  gravity  increases  with  the  atomic  weight,  as  does 
also  the  atomic  volume;  on  the  other  hand,  there  is  a  fall  in  the 
melting-  and  boiling-points. 

From  a  chemical  standpoint  we  notice,  in  the  first  place,  the 
same  general  type  in  the  compounds,  showing  that  all  these  ele- 
ments are  univalent.  The  hydroxides  all  have  the  formula  ROH. 
the  halogen  compounds  RX,  etc.  The  salts  of  them  all,  even 
the  carbonates  and  phosphates,  are  soluble  in  water  (although  in 
different  degrees),  the  carbonates  with  basic  reaction.  The  metals 
all  oxidize  very  readily  in  the  air. 

On  the  other  hand,  we  cannot  overlook  the  fact  that  the  metals 
potassium,  rubidium  and  cesium,  which  are  very  similar  to  each 
other,  differ  from  sodium  and  lithium  in  many  respects.  The  last- 
named  metal,  as  we  shall  see  in  the  sequel,  displays  analogy 
with  magnesium  in  several  important  points,  thus  differing  from 
the  metals  of  its  own  group.  A  slight  divergence  in  the  behavior 
of  the  first  members  of  a  group  from  that  of  the  rest  is  found  to 
characterize  almost  all  of  the  groups.  We  may  recall  carbon,  for 
instance,  the  first  member  of  the  fourth  group,  which  differs  dis- 
tinctly from  silicon  and  the  rest  in  the  ability  of  its  atoms  to  unite 
with  each  other;  also  fluorine  with  its  soluble  silver  compound. 
Still  other  examples  of  this  sort  will  be  met  with  later. 

Sodium  differs  from  the  sub-group,  K,  Rb,  Cs,  in  the  solubility 
of  its  salts.  The  sodium  salts  are  almost  all  readily  soluble  in 
water;  this  is  true  even  of  the  platinum  double-salt,  Na2PtCl6,  the 
acid  sodium  tartrate  and  others.  Soda  crystals  effloresce,  while 
potash  deliquesces  in  the  air.  The  spectra  of  sodium,  on  the  one 
hand,  and  the  other  alkali  metals,  on  the  other,  are  entirely  dis- 
similar. 

Ammonium  Salts. 

234.  In  the  description  of  ammonia  (§  112)  it  was  already 
observed  that  it  combines  with  acids  directly,  forming  salts  which 
are  very  similar  to  those  of  potassium,  and  in  which  the  group, 
NH4,  the  ammonium  group,  is  assumed  to  exist.  In  connection 
with  the  alkali  group  a  description  of  a  few  ammonium  salts  may 
find  a  place. 

The  aqueous  solution  of  ammonia  must,  because  of  its  electrical 
conductivity  and  its  alkaline  reaction,  contain  NH4  and  OH'  ions 


332  INORGANIC  CHEMISTRY.  [§234- 

and  hence  also  undissociated  molecules  of  ammonium  hydroxide, 
NH4OH.  While  solutions  of  the  alkalies,  KOH,  NaOH,  etc., 
conduct  the  electric  current  very  well,  this  is  not  the  case  with  an 
ammonia  solution;  it  is  a  poor  conductor.  A  0.1  normal  solution 
contains  only  5%  of  ionized  NH4OH  molecules,  while  a  solution 
of  potassium  hydroxide  of  the  same  concentration  is  91% 
ionized. 

An  aqueous  solution  of  ammonia  may  be  presumed  to  contain: 
(1)  free  ammonia,  NH3;  (2)  hydrates  of  ammonia,  NH3-n  aq.;  (3) 
ammonium  hydroxide,  NH4OH;  (4)  the  ions  NH4  and  OH'.  The 
existence  of  these  hydrates  in  addition  to  free  ammonia  reveals 
itself  in  the  behavior  of  ammonia  solutions  on  being  shaken  with 
chloroform.  According  to  BERTHELOT'S  law  (ORG.  CHEM.,  §  25) 
the  distribution  ratio  of  the  ammonia  between  the  two  solvents 
should  be  constant.  But  this  is  not  the  case.  Therefore,  just 
as  the  deviations  from  HENRY'S  law  lead  us  to  conclude  that  a 
dissolved  gas  exists  in  a  special  condition,  so  we  can  apply  a  simi- 
lar explanation  to  the  exceptions  to  the  BERTHELOT  law;  for 
HENRY'S  law  is  really  the  expression  of  the  distribution  ratio  of 
a  gas  between  a  liquid  and  a  vacuum,  while  the  other  law  has  to 
deal  with  the  distribution  ratio  between  two  liquids. 

Since  in  the  case  of  ammonia  this  deviation  is  observed  only 
when  one  of  the  two  liquids  is  water,  we  are  obliged  to  conclude 
that  there  is  a  combination  of  the  ammonia  with  the  water.  The 
reason  for  assuming  the  existence  of  hydrates  instead  of  ammo- 
nium hydroxide,  NH4OH,  is  a  double  one.  We  have,  on  the 
one  hand,  the  analogy  between  the  behavior  of  ammonia  and 
amines  and,  on  the  other,  the  entirely  abnormal  behavior  of  the 
organic  quaternary  ammonium  bases.  For,  while  the  aqueous 
solutions  of  primary,  secondary  and  tertiary  amines  are  weak 
electrolytes,  as  is  the  case  with  ammonia,  the  solution  of  a  quater- 
nary base,  on  the  contrary,  conducts  electricity  as  well  as  a  solu- 
tion of  potassium  or  sodium  hydroxide.  We  may  thus  conclude 
that  if  ammonium  hydroxide,  NH4OH,  could  reach  the  same 
concentration  in  solution  as  a  quaternary  base  it  would  display 
just  as  great  a  conductivity  as  the  latter.  Unlike  the  quaternary 
base,  however,  it  breaks  up  principally  into  ammonia,  NH3,  and 
water,  for  the  quaternary  base  cannot  be  thus  decomposed. 

That  an  aqueous  ammonia  solution  really  contains  at  least 


§234.]  AMMONIUM  SALTS.  333 

an  appreciable  quantity  of  the  hydroxide  is  evidenced  by  the 
existence  of  its  ions.  These  necessitate  the  establishment  of  the 
equilibrium 


however  far  the  point  of  equilibrium  may  be  displaced  toward 
the  right. 

The  great  tendency  of  ammonium  hydroxide  to  break  up  into 
ammonia  and  water  is  the  reason  for  the  very  feeble  basic  reac- 
tion of  an  aqueous  solution  of  ammonia,  for  undoubtedly  ammo- 
nium hydroxide,  so  far  as  it  is  formed,  is  extensively  ionized,  like 
the  strong  bases.  This  view  is  supported  by  various  observa- 
tions, among  them  the  neutral  reaction  of  the  ammonium  salts 
of  strong  acids,  such  as  the  chloride  and  the  nitrate,  and  also 
the  alkaline  reaction  of  the  carbonate  and  the  cyanide,  in  har- 
mony with  the  similar  alkaline  reaction  of  the  corresponding  salts 
of  the  alkali  metals. 

Ammonium  chloride,  NH^Cl,  sal  ammoniac,  is  obtained  from 
the  ammonia  liquor  of  the  gas  factories  (§  112).  The  ammonia 
is  expelled  by  warming  and  absorbed  in  hydrochloric  acid;  this 
solution  is  evaporated  and  the  solid  residue  sublimed,  whereby 
the  salt  is  obtained  in  compact  fibrous  masses.  It  dissolves  in 
2.7  parts  of  cold,  and  in  1  part  of  boiling,  water  and  crystallizes 
out  of  the  solution  in  small,  usually  feather-like  groups  of  octa- 
hedrons or  cubes.  It  has  a  sharp  saline  taste. 

Ammonium  chloride  vaporizes  easily,  dissociating  into  ammonia 
and  hydrochloric  acid,  as  is  shown  by  the  vapor  density,  which  at 
350°  is  only  half  as  great  as  calculated. 

This  dissociation  can  be  easily  demonstrated  in  the  following  manner  : 
Introduce  into  a  tube  sealed  at  one  end  a  little  ammonium  chloride  and, 
not  far  from  this,  a  piece  of  blue  litmus  paper.  In  front  of  the  latter  is 
pushed  a  plug  of  asbestos  wool  and  finally  a  piece  of  red  litmus  paper. 
The  chloride  is  then  heated.  Since  hydrochloric  acid  has  a  smaller 
diffusion  velocity  than  ammonia  the  latter  passes  through  the  wad 
first  and  colors  the  red  paper  blue  ;  as  a  result  an  excess  of  hydrochloric 
acid  is  left  at  the  other  end  and  it  reddens  the  blue  paper  placed  there. 

It  is  a  remarkable  fact,  discovered  by  BAKER,  that  perfectly  dry 
ammonium  chloride  (having  stood  for  a  long  time  in  a  desiccator 
over  resublimed  phosphorus  pentoxide)  has  the  normal  vapor 
density.  On  the  other  hand,  the  same  investigator  found  that 


334  INORGANIC  CHEMISTRY.  [§§234, 

similarly  dried  ammonia  gas  and  hydrochloric  acid  gas  do  not 
unite  to  form  ammonium  chloride  (§  38).  Traces  of  water  thus 
produce  a  marked  catalytical  acceleration,  both  of  the  formation 
and  of  the  decomposition  of  ammonium  chloride.  We  have  here 
an  illustration  of  the  general  rule  that  when  one  part  of  the 
system  in  a  reversible  reaction  is  accelerated  by  a  catalyzer  the 
other  must  be  likewise  affected.  The  proof  of  this  rule  lies  in 
the  impossibility  of  the  contrary  being  true,  since  that  would 
necessitate  a  change  in  the  equilibrium  (see  §  49). 

In  many  other  cases  it  is  also  observed  that  traces  of  water  have  a 
considerable  influence  on  the  velocity  of  chemical  reactions.  The  follow- 
ing examples  may  be  cited:  (1)  Phosphorus,  that  ordinarily  takes  fire 
in  moist  air  at  a  little  above  room  temperature,  can  be  heated  in  oxygen 
to  150°  without  ignition,  provided  the  oxygen  has  been  carefully  dried  by 
phosphorus  pentoxide.  (2)  Carbon  monoxide  burns  in  moist  oxygen  much 
more  easily  than  in  dry  oxygen.  (3)  Very  carefully  dried  detonating- 
gas  can  be  heated  in  a  tube  to  red-heat  without  exploding. 

Ammonium  sulphate,  (NH4)2S04,  crystallizes  in  large  rhombic 
prisms  and  dissolves  very  readily  in  water.  On  boiling  the  aque- 
ous solution  some  ammonia  escapes,  acid  sulphate  being  formed. 

Its  solution  in  30%  hydrogen  peroxide  yields  on  evaporation  crystals 
of  the  composition  (NH4)2SO4-H2O2.  When  these  are  heated  under  re- 
duced pressure  a  high  per  cent  hydrogen  peroxide  distils  off. 

Ammonium  nitrate,  NH4N03,  deliquesces  in  the  air;  when 
heated  it  breaks  up  into  water  and  nitrous  oxide  (§  119).  This 
salt  is  known  in  three  crystallized  modifications,  the  transition 
points  (§  70)  of  which  have  been  determined. 

Ammonium  phosphates. — The  tertiary  salt,  (NH4)3PO4,  is 
deposited  in  crystalline  form  on  mixing  concentrated  solutions  of 
phosphoric  acid  and  ammonia.  It  cannot  be  dried,  however,  for 
it  then  loses  ammonia  and  goes  over  into  the  secondary  phosphate, 
(NH4)2HPO4.  On  boiling  the  solution  the  salt  again  yields  ammo- 
nia and  is  transformed  into  the  primary  phosphate. 

The  best  known  of  these  salts  is  the  sodium  ammonium  phos- 
phate, NaNH4HP04  •  4H2O,  microcosmic  salt.  It  forms  large 
transparent  crystals.  On  being  heated  it  melts,  loses  water  and 
ammonia,  and  passes  over  into  a  vitreous  substance,  sodium  meta- 
phosphate,  NaPO3. 

Ammonium  carbonate  was  formerly  obtained  by  the  dry  distil- 
lation of  nitrogenous  organic  substances,  such  as  hair,  nails, 
leather,  etc.,  hence  the  name  "  salt  of  hartshorn,"  which  still 


235.]  SALT  SOLUTIONS.  335 

clings  to  it.  At  present,  however,  it  is  made  by  dry  distilling  a 
mixture  of  calcium  carbonate  and  ammonium  chloride  or  sulphate. 
The  product  is  a  mixture  (molecule  for  molecule)  of  acid  salt, 
NH4HCOs,  and  ammonium  carbamate,  NH2-CO2-NH4  (this 
latter  being  the  neutral  salt  minus  1H2O).  From  its  composition, 
(NH3)3(CO2)2-H20,  it  takes  the  name  ammonium  sesquicarbonate, 
On  passing  ammonia  gas  into  a  concentrated  aqueous  solution  of  it 
the  neutral  salt,  (NH4)2CO3,  separates  out  as  a  crystalline  powder; 
it  smells  strongly  of  ammonia  and  passes  slowly  over  into  the  acid 
salt,  NH4HCO3,  a  white  odorless  powder,  which  is  scarcely  soluble 
in  water.  This  acid  salt  is  also  formed  directly  from  the  sesqui- 
carbonate, as  the  latter  gives  off  carbon  dioxide  and  ammonia  in 
the  air  (hence  the  odor  of  ammonia)  and  goes  over  into  the  first- 
named  salt. 

Ammonium  sulphide  is  extensively  used  in  analysis  (§  73).  A 
solution  of  ammonium  hydrosulphide  (or  sulphydrate),  NH4SH, 
is  obtained  by  saturating  aqueous  ammonia  with  hydrogen  sulphide; 
it  is  a  colorless  liquid,  which  soon  turns  yellow  because  of  the  for- 
mation of  ammonium  polysulphides.  The  oxygen  of  the  air  oxidizes 
part  of  the  hydrogen  sulphide  and  thus  sets  free  sulphur,  which 
combines  with  ammonium  hydrosulphide  to  form  polysulphides. 
These  polysulphides  are  also  obtained  by  dissolving  sulphur  in  a 
solution  of  ammonium  hydrosulphide. 

On  mixing  2  vols.  NH3  gas  and  1  vol.  H2S  gas  at  —  18°  a  white  crystal- 
line mass  is  obtained,  which  decomposes  at  ordinary  temperatures  into 
NH4SH  and  NH3.  The  compound  NH4SH  separates  out  crystalline 
when  hydrogen  sulphide  is  passed  into  alcoholic  ammonia.  As  low  as 
45°  it  is  completely  dissociated  into  equal  volumes  of  NH3  and  H2S. 

SALT  SOLUTIONS. 

235.  Every  solid  substance  is  soluble  in  every  liquid;  however, 
the  proportion  which  dissolves  can  vary  all  the  way  from  zero  to 
infinity.  If  only  an  infinitesimal  amount  of  the  solid  goes  into 
solution,  we  say  ordinarily  that  the  substance  is  "  insoluble  "  in  the 
liquid ;  there  can  be  no  doubt,  however,  that,  if  our  means  of  inves- 
tigation were  sufficiently  improved  and  large  enough  quantities  of 
Jiquid  were  taken,  the  solubility  would  be  perceptible.  This  has 
already  been  demonstrated  in  many  cases  of  so-called  insoluble 
substances  (§  210).  Even  when  we  confine  our  attention  to 
aqueous  solutions  of  salts  (including  acids  and  bases)  we  find  the 


336  INORGANIC  CHEMISTRY.  [§§235- 

same  infinite  difference  in  solubility  that  is  observed  between  sub- 
stances in  general.  Substances  such  as  sand,  barium  sulphate 
(§  262),  silver  iodide,  etc.,  are  "  insoluble  ";  others,  like  sulphuric 
acid,  are  able  to  dissolve  in  any  given  amount  of  water. 

With  regard  to  the  solubility  of  salts  the  following  practical  rules  are 
worth  remembering :  Potassium,  sodium  and  ammonium 
salts  are  soluble. — Normal  nitrates,  chlorates  and  ace- 
tates are  soluble.— Normal  chlorides  are  soluble  (except  AgCl, 
Hg2Cl2,  and  PbCl2).— Normal  s  ulp  h  a  t  e  s  are  soluble  (except  those 
of  Ba,  Sr,  Ca,  and  Pb) . — H  ydroxides  are  insoluble  (except  those 
of  the  alkalies  and  alkaline  earths) . — Normal  carbonates,  phos- 
phates, and  sulphides  are  insoluble  (except  those  of  the  alka- 
lies).— Basic  salts  are  insoluble. — Acid  salts  are  soluble  if 
the  acid  iteelf  is  soluble. 

The  solubility,  i.e.  the  maximum  relative  amount  of  salt  that 
can  go  into  solution,  is  a  function  of  the  temperature  and  the 
pressure.  In  the  great  majority  of  cases  the  solubility  increases 
with  the  temperature.  If  the  temperature  is  plotted  on  the  axis 
of  abscissas  and  the  amount  of  salt  which  dissolves  in  one  hundred 
parts  of  water  is  plotted  on  the  ordinate  axis,  a  solubility  curve 
is  obtained  (Fig.  52)  which  shows  at  a  glance  the  variation  of  the 
solubility  with  the  temperature. 

For  some  salts,  e.g.  potassium  nitrate,  the  solubility  increases 
very  rapidly  with  the  temperature;"  for  sodium  chloride  it  remains 
practically  constant.  In  certain  cases,  such  as  those  of  calcium 
hydroxide  and  calcium  sulphate  (within  certain  limits  of  tempera- 
ture) the  solubility  decreases  with  rising  temperature.  These 
phenomena  are  connected,  as  has  already  been  explained,  with 
the  heat  of  solution,  i.e.  with  the  caloric  effect  which  accompanies 
the  process  of  solution,  and  in  the  manner  expressed  by  VAN'T 
HOFF'S  principle  of  mobile  equilibrium  (§  103).  In  fact  saltpetre, 
for  instance,  whose  solubility  increases  very  rapidly  with  the 
temperature  (see  Fig.  52)  dissolves  in  water  with  a  considerable 
absorption  of  heat.  $. 

236.  The  term  heat  of  solution  has  various  meanings.  We  are  obliged 
to  distinguish  between  (1)  the  caloric  effect  of  dissolving  a  salt  in  a  very 
large  amount  of  water;  (2)  the  caloric  effect  of  dissolving  a  salt  in  an 
almost  saturated  solution;  and  (3)  the  total  heat  of  solution,  i.e.  the 
whole  caloric  effect  of  dissolving  a  salt  in  water  until  the  solution  is 
saturated.  As  a  rule  these  three  magnitudes  will  have  dissimilar 
values,  indeed  their  algebraic  signs  may  be  opposite.  This  is  the  case, 


236.] 


SALT  SOLUTIONS. 


337 


for  instance,  with  the  compound  CuCl2-2H20;  1  g.-mol.  dissolved  fa 
198  g.-mols.  H2O  at  11°  gives  a  caloric  effect  of  +3.71  Cal.;  19.56 
g.-mols.  in  the  same  amount  of  water,  —3.129  Cal. 

The  heat  of  solution  to  which    VAN'T  HOFF'S  principle  applies  is 
that  of  the  salt  in  its  saturated  solution.     We  have  here  the  system: 


10"          20U          30°         *«°          50°          60°          70C 

FIG.  52. — SOLUBILITY  CURVES. 


100° 


salt  M-  saturated  solution;  when  the  temperature  changes,  the  equi- 
librium is  displaced,  i.e.  salt  either  goes  into  solution  or  crystallizes  out, 
the  latter  action  producing  just  as  large  a  thermal  effect  numerically 
as  dissolving  in  the  saturated  solution,  but  with  the  opposite  sign.  Since 
this  was  not  taken  into  consideration  when  the  matter  was  first  dis- 


338  INORGANIC   CHEMISTRY.  [§§236- 

cussed,  it  was  believed  that  there  were  exceptions  to  the  principle,  but 
closer  investigation  has  proved  the  contrary. 

In  some  cases  the  solubility  of  a  salt  at  first  increases  gradually 
with  rising  temperature  and  then  steadily  decreases,  so  that  the 
solubility  curve  has  a  maximum  (cf.  Fig.  53).  In  full  agreement 
with  VAN'T  HOFF'S  principle  the  heat  of  solution  is  negative  in 
the  ascending  portion  of  the  curve,  zero  at  the  maximum  and 
positive  in  the  descending  portion.  In  the  case  of  gypsum, 
CaSO4-2H2O,  for  instance,  the  maximum  was  found  to  lie  at 
about  38°  and  at  that  point  the  heat  of  solution  was  actually 
proved  to  be  0.00;  at  14°  it  is  -0.36;  above 
35°,  +0.24. 

The  effect  of  pressure  on  the  solubility 
is  at  the  most  very  slight,  but  it  is  in  entire 
accord  with  the  principle  of  LE  CHATELIER. 
^  Ammonium  chloride,  for  instance,  dissolves 

with  expansion;  therefore  its  solubility  lessens 
t  IG.  53.  .  , .  „_,  .  . 

with  increasing  pressure  (1%  for  an  increase 

of  160  atm.).  Copper  sulphate,  which  dissolves  with  contraction, 
has  its  solubility  increased  3.2%  by  an  increase  of  60  atm.  pressure. 

237.  It  was  formerly  thought  that  the  terms  "  solvent  "  and 
"  dissolved  substance  "  ("  solute  ")  should  be  kept  distinct.  How- 
ever, it  has  since  developed  that  there  is  no  essential  difference 
between  the  components  of  a  solution,  and  that  aqueous  solutions 
are  therefore  better  defined  as  "  liquid  complexes,  one  of  whose 
components  is  water,"  than  as  "  water  in  which  substances  are 
dissolved." 

The  interchangeability  of  the  terms  "  solvent "  and  "  solute  " 
is  evidenced  first  of  all  by  the  phenomena  attending  the  cooling 
of  salt  solutions.  Let  us  consider,  for  instance,  a  nearly  saturated 
solution  of  potassium  chloride  at  a  definite  temperature.  .  We  have 
in  it  two  substances  (KC1  and  H20)  and  two  phases  (§  p.),  hence 
two  degrees  of  freedom.  We  will  suppose  that  the  solution  is  then 
cooled;  potassium  chloride  crystallizes  out  forthwith  and,  as  three 
phases  are  then  present,  the  system  becomes  univariant  We  recall 
that  changes  in  the  quantity  of  any  phase  have  no  effect  on  such 
a  system;  therefore,  if  more  salt  is  introduced  into  the  system, 
the  concentrations  of  saturated  solution  and  vapor  are  unaffected. 


237.]  SALT  SOLUTIONS.  339 

This  is  none  the  less  true  when  water    is  added  or  the  vapor 
volume  increased,  so  long  as  the  three  phases  remain. 

On  cooling  still  farther,  more  potassium  chloride  is  gradually 
deposited  until  a  point  is  reached  below  which  the  entire  liquid 
congeals  to  a  mixture  of  salt  and  ice.  This  point  is  known  as  the 
cryohydric,  or  eutectic,  point.  There  are  now  four  phases, — salt, 
ice,  solution  and  vapor; — hence  the  system  has  become  nonvariant. 

The  opinion  was  formerly  held  that  at  this  point  a  chemical  com- 
pound between  the  salt  and  water  (a  ' '  cryohydrate  ")  came  into  exist- 
ence. That  it  is  only  a  matter  of  mixtures  can  be  seen  in  the  case  of 
colored  salts  (K2OO4),  for  instance,  with  a  microscope;  moreover  the 
composition  of  these  so-called  hydrates  may  differ  in  case  the  solidifi- 
cation takes  place  under  a  different  pressure. 

If  we  start  with  a  dilute  potassium  chloride  solution  as  another 
example,  and  cool  it,  we  have  ice  formed  at  a  definite  temperature 
and  a  univariant  system  established,  ice  being  the  third  phase 
required.  Below  this  point  the  solution  can  be  regarded  as  saturated 
in  respect  to  ice,  just  as  it  could  be  considered  saturated  in 
respect  to  the  salt  in  the  previous  case;  for  an  increase  of  the 
solid  phase  (ice)  does  not  now  cause  a  displacement  of  the 
equilibrium  (§  71)  any  more  than  the  addition  of  the  solid  (salt) 
did  in  the  previous  instance.  The  addition  of  potassium  chloride 
causes  part  of  the  ice  to  go  into  solution  (i.e.  melt) ;  for  the 
dissolving  of  more  salt  increases  the  concentration  of  the  solution. 
Therefore,  if  the  temperature  is  kept  constant,  ice  must  melt  in 
order  to  restore  the  solution  to  its  previous  concentration.  It  is 
therefore  perfectly  analogous  to  the  addition  of  water  to  a  satu- 
rated potassium  chloride  solution  in  contact  with  the  solid  salt, 
in  which  case  also  the  solid  phase  goes  into  solution.  If  the  tem- 
perature rises,  more  ice  dissolves;  if  it  falls,  more  crystallizes  out — 
just  as  with  rising  temperature  more  potassium  chloride  goes  into 
solution  .^nd  with  sinking  temperature  more  crystallizes  out.  On 
farther  cooling  more  and  more  ice  will  be  deposited  until,  in  this 
case  also,  the  cryohydric  point  is  reached,  below  which  the  whole 
system  solidifies  to  a  mixture  of  salt  and  ice.  The  analogy  is 
therefore  complete. 

The  cryohydric  point  is,  according  to  this  view,  the  point  of 
intersection  of  two  curves,  viz.:  the  solubility  curves  of  salt  and 
of  ice  in  the  salt  solution. 


340  INORGANIC  CHEMISTRY.  [§  237. 

Another  argument  against  the  assumption  of  any  essential 
difference  between  solvent  and  solute  is  found  in  the  behavior  of 
the  solutions  of  certain  hydrous  salts,  e.g.  CaCl2-6H20.  A  satu- 
rated solution  of  CaCl2  in  water  at  30.2°  has  exactly  the  com- 
position CaQ2-6H20.  At  this  temperature,  therefore,  the  hydrate 
melts  to  a  homogeneous  liquid.  //  either  H20  or  CaCl2  is 
added,  there  is  a  deposition  of  CaCl2  -  6H20  on  cooling,  for  the  addi- 
tion of  either  causes  a  depression  of  the  point  of  solidification  (freez- 
ing-point) of  CaCl2-6H20.  In  the  first  case  this  hydrate  is  in 
equilibrium  with  a  liquid  which  contains  more  water  than  the 
hydrate  does  and  which  is  therefore  called  an  aqueous  solution 
in  the  ordinary  sense.  In  the  second  it  is  in  equilibrium  with  a 
liquid  which  contains  more  CaCl2  than  CaCl2-6H20  and  must 
therefore  be  regarded  as  a  solution  in  CaCl2. 

On  examining  the  solubility  curves  of  various  salts  (cf.  Fig.  52) 
it  is  found  that  they  are  in  general  regular;  however,  in  one  of  the 
curves  (sodium  sulphate)  a  sudden  change  of  direction  is  noticed. 
This  is  often  observed  with  salts  that  contain  water  of  crystalliza- 
tion. Taking  sodium  sulphate  as  an  example,  the  phenomenon  may 
be  explained  thus:  It  has  already  been  remarked  (§  225)  that  this 
salt  has  a  transition  point  at  the  temperature  of  33°,  Na2S04-  10H2O 
being  transformed  into  Na2S04  and  10H20.  Up  to  33°,  therefore, 
we  have  the  hydrous  salt  as  the  solid  phase;  above  this  temperature 
the  anhydrous  salt.  This  change  must  necessarily  involve  a  sudden 
bend  of  the  solubility  curve.  Below  33°  the  curve  represents  the 
solubility  of  Na2SO4-10H2O,  above  33°  that  of  Na2S04.  We 
can  therefore  also  regard  the  point  of  inflection  of  the  curve  (at  33Q) 
as  the  point  of  intersection  of  the  curves  for  Na2S04-10H20  and 
Na2S04.  In  sodium  sulphate  the  special  case  appears  where  the  sol- 
ubility of  the  anhydrous  salt  decreases  with  rising  temperature  and 
hence  the  solubility  curve  falls  as  the  temperature  rises  above  33°. 

In  the  light  of  the  above  the  solubility  of  a  substance  which 
has  a  transition  point  is  the  same  for  both  modifications  at  this 
point.  This  must  always  be  the  case;  it  can  be  demonstrated  in  the 
same  way  as  in  §  70,  where  it  was  shown  that  the  vapor  pressures 
become  equal  at  the  transition  point.  Indeed  the  same  figure  can 
be  employed,  if  it  is  borne  in  mind  that  the  solubility  of  a 
metastable  modification  is  always  greater  than  that  of  the  stable 
modification  at  one  and  the  same  temperature.  Inversely,  more- 
over, we  have  here  a  means  of  determining  the  transition  point. 


§  237-1 


SALT  SOLUTIONS. 


341 


In  general,  as  OSTWALD  has  pointed  out,  the  solubility  of  any 
substance  whatever  is  dependent  on  the  condition  in  which  it  exists. 
The  solid  phase  determines  the  equilibrium,  not  only  in  virtue  of 
its  chemical  composition  but  also  by  the  particular  modification 
in  which  the  solid  substance  is  present.  Thus,  e.g.,  each  of  the 
various  forms  of  the  same  polymorphous  substance  or  different 
hydrates  of  the  same  salt  has  its  own  solubility,  other  things  being 
equal. 

In  a  hydrous  salt  we  may  have  the  case  where  there  are  various 
hydrates,  which  are  connected  with  each  other  by  transition  points. 
A  salt  with  m  +  n  molecules  of  water  of  crystallization  passes  over 
at  a  definite  temperature  into  another  with  m  molecules,  for  ex- 
ample. The  latter  may,  at  a  higher  temperature,  have  a  second 
transition  point  (to  anhydrous  salt).  At  each  of  these  points  the 
solubility  curve  will  show  a  bend,  because  the  solid  phase  changes; 
the  curve  will  therefore  assume  some  such  form  as  t.^at  of  Fig.  54. 

Let  us  examine  such  a  solubility  curve  a  little  more  closely. 
At  0°  (A  in  F.g.  rj4)  we  will  suppose  that  we  have  pure  water  and 
ice,  to  start  with,  and  JQO 
that  small  portions  of 
salt  are  then  gradually  1 
dissolved.  If  the  ice  Eg 
phase  is  to  be  preserved,  2  o 
the  temperature  must  £ 
be  allowed  to  sink,  fora  2 
salt  solution  has  a  lower 
freezing-point  than  pure 
water.  We  therefore 


TEMPERATURE 


FIG.  54. 


move  along  the  curve  AK.  Soon  a  point  K  is  reached  when  no 
more  salt  dissolves,  since  all  the  water  has  now  turned  to  ice. 
Here,  therefore,  we  have  a  mixture  of  ice  and  solid  salt,  or,  in 
other  words,  the  cryohydric  point. 

If  we  wish  to  bring  more  salt  into  solution  after  K  is  reached 
the  temperature  must  be  raised.  The  ice  phase  then,  of  course, 
disappears  and  in  its  place  we  have  the  salt  with  m+n  molecules 
of  water  of  crystallization  as  solid  phase.  If  the  temperature  is 
steadily  raised  and  the  solution  is  kept  constantly  saturated  by 
the  addition  of  this  salt,  we  move  along  the  curve  KB.  At  B, 
however,  we  meet  the  transition  point  from  the  salt  with  m+n 
mols.  H2O  to  the  one  with  m  mols.  H^O;  hence  the  solubility 


342  INORGANIC  CHEMISTRY.  [§  237. 

curve  must  again  bend  here  and  in  such  a  way  that  atthepoint 
B  the  solubility  curve  of  the  salt  with  m  +  n  mols.  H2O  is  steeper 
than  that  of  the  salt  with  m  mols.,  no  matter  what  the  form  of 
the  curves  KB  and  BC  may  be.  This  is  readily  understood  by 
a  course  of  reasoning  entirely  analogous  to  that  given  for  the 
transition  of  ice  to  water  or  of  rhombic  to  monoclinic  sulphur 
(§  70).  Finally  at  C  we  have  a  second  transition  point  from  salt 
with  m  mols.  H^O  to  anhydrous  salt,  so  that  the  solubility  curve 
there  shows  one  more  bend.  Where  the  curve  CD  ends  depends 
on  circumstances.  In  many  cases,  e.g.  that  of  silver  nitrate,  it  ends 
at  the  melting-point  of  the  anhydrous  salt  (concentration  of  the 
solution  =  100%).  In  other  instances  the  anhydrous  salt  can  form 
a  second  (fused)  liquid  layer  under  the  saturated  solution.  Finally, 
mention  may  also  be  made  of  the  case  of  copper  sulphate,  which  at 
a  given  temperature  loses  its  water  of  crystallization  in  contact  with 
its  saturated  solution  and  from  that  point  .on  shows  a  decrease  in 
solubility  with  rising  temperature,  which  finally  ends  in  almost 
total  insolubility. 

If  we  draw  a  line  kik2  through  K  parallel  to  the  ordinate  axis, 
the  figure  is  divided  by  this  line  and  AKBCD  into  the  following 
regions:  To  the  right  of  the  solubility  curve  is  the  region  of  the 
unsaturated  solutions,  AKk2  is  that  of  the  superfused,  k±KBCD 
that  of  the  supersaturated,  solutions.  To  the  left  of  k\k2  only 
ice  +  solid  salt  can  exist  under  ordinary  pressures. 

A  solubility  curve,  such  as  that  represented  in  Fig.  54,  can, 
on  the  other  hand,  be  used  to  detect  the  existence  of  compounds 
between  the  salt  and  the  water.  From  the  cryohydric  point 
upward  every  bend  in  the  solubility  curve  shows  that  a  salt  with 
a  different  amount  of  water  of  crystallization  has  been  formed. 
Each  branch  of  the  curve  thus  represents  a  separate  salt,  i.e.  a 
different  solid  phase.  The  composition  of  these  solid  phases  is 
by  no  means  always  self-evident.  Such  is,  however,  the  case  when 
the  phase  fuses  without  altering  its  composition,  or,  what  amounts 
to  the  same  thing,  when  it  can  exist  in  equilibrium  with  a  liquid 
phase  of  the  same  composition.  This  does  not  often  occur  with 
salts  in  aqueous  solution,  but  an  example  of  it  was  seen  above  in 
the  case  of  CaCl2-6H20.  The  inspection  of  the  solubility  curve 
or  melting-point  curve  is  then  especially  valuable  for  the  discovery 
of  compounds.  In  order  to  understand  this  let  us  first  examine 
a  system  of  two  substances,  A  and  B,  which  do  not  combine.  Fig. 


§  237.] 


SALT  SOLUTIONS. 


343 


55,  represents  the  melting-point  curve  that  one  obtains  on  the 
addition  of  increasing  amounts  of  B  to  A.  At  first  the  melting- 
point  sinks  until  the  eutectic  point  *  E  is  reached.  Along  AE  A 
alone  separates  out  of  the  fused  mass  on  freezing.  At  E,  however, 
B  also  separates  out.  If  more  of  B  is  now  added  the  melting- 
point  rises;  we  obtain  the  curve  EB,  which  terminates  in  the 
melting-point  of  pure  B.  Along  EB  only  B  separates  out  of  the 
fused  mass. 


100 


FIG.  55. 


20    so   IB   w 
FIG.  56. 


100 


Suppose  we  now  assume  that  A  and  B  form  a  compound  AB 
in  the  molecular  proportions  1:1  (Fig.  56).  On  the  addition  of 
B  to  A  AB  is  formed  and  dissolves  in  the  excess  of  A.  This  lowers 
its  melting-point.  When  a  certain  amount  of  B  has  been  added 
this  point  is  lowered  to  E\.  Here  both  A  and  AB  separate  out. 
EI  is  the  eutectic  point  for  mixtures  of  A  and  AB.  If  more  of 
B  is  added  the  melting-point  rises,  just  as  in  the  case  where 
there  is  no  combination  between  the  components.  Only  AB 
now  separates  out  of  the  fused  mass.  The  continued  addition 
of  B,  however,  increases  the  amount  of  the  compound  AB;  at 
M  free  A  has  disappeared  and  the  mass  consists  wholly  of  pure 
AB,  whose  melting-point  is  M.  At  this  point  the  melting-tem- 
perature reaches  its  maximum,  for  the  addition  of  either  A  or  B 
lowers  the  melting-point  of  the  pure  compound.  The  further 

*  The  term  "eutectic  point"  is  more  general  than  "cryohydric  point,"  the 
latter  term  being  usually  restricted  to  aqueous  solutions.  "Eutectic  mix- 
ture" and  "cryohydric  mixture"  ("cryohydrate")  are  similarly  related. 


344 


INORGANIC  CHEMISTRY. 


[§237 


course  of  the  curve  is  readily  seen.  As  more  and  more  B  is  added 
to  AB  the  melting-point  sinks,  AB  alone  separating  out,  until 
the  eutectic  point  E2  is  reached,  where  both  AB  and  B  crystallize 
out,  and  thereafter  the  melting-point  again  rises  along  E2B  till 
it  finally  ends  in  the  melting-point  of  the  pure  substance  B. 

If  more  than  one  compound  is  formed  between  A  and  B,  each 
one  will  cause  a  maximum  point  in  the  curve,  i.e.,  each  maximum 
will  correspond  to  a  compound.  The  following  examples  will  serve 
to  make  this  clear: 

1.  The  system  S03+H20.  Here  there  are  a  number  of  hy- 
drates, which  are  indicated  by  the  melting-point  curve  (Fig.  57). 


/\ 


\       / 


1 62     05 


-35 


FIG.  57. 

A  mixture  of  62%  SO3  +  38%  H2O  has  a  freezing-point  of  —20°. 
On  the  addition  of  sulphur  trioxide  this  point  rises  till  at  the  com- 
position SO3  +  2H2O=H2SO4-H2O  it  reaches  its  first  maximum. 
At  this  temperature  (8°)  the  whole  mass  solidifies,  yielding  crystals 
of  the  above  composition.  The  further  addition  of  sulphur  tri- 
oxide has  the  same  effect  on  the  melting-point  of  the  hydrate 
H2S04-H2O  as  the  ordinary  addition  of  a  foreign  substance  to 
a  pure  substance.  At  first  the  melting-point  falls;  that  which 
crystallizes  out  is  the  hydrate  H2S04-H20.  At  a  composition 
of  about  3H2O  +  2SO3  a  eutectic  point  is  reached  (§237).  The 
mass  solidifies  at  —35°  to  a  mixture  of  H2S04-H2O  and  H2S04. 
Continued  addition  of  sulphur  trioxide  causes  a  rise  of  the  melting- 
point  till  at  +10°  a  second  maximum  is  reached,  where  the  whole 
solidifies  to  a  homogeneous  mass,  consisting  of  pure  sulphuric 


§  237.] 


SALT  SOLUTIONS. 


345 


acid,  H2SC>4.  Along  this  ascending  branch  of  the  curve  H2S04 
crystallizes  out.  The  melting-point  curve  then  proceeds  to  a 
third  maximum  point,  corresponding  to  the  hydrate  H^SC^+SOs 
=  H2S2O7  (pyrosulphuric  acid),  and  comes  to  an  end  in  the  melting- 
point  of  the  asbestine  form  of  sulphur  trioxide  at  +40°. 

2.  In  non-aqueous  liquids  the  relations  are  almost  exactly 
the  same,  as  may  be  seen  from  a  consideration  of  the  system 
S  +  C1.  It  was  remarked  in  §  75  that,  while  the  compound  SCU 
could  not  be  isolated,  the  form  of  the  melting-point  curve  left  nc 


100      90       80       70       60       50       40       30       20       10 
S  Atomic  Percentage  of  Sulphur 

FIG.   58. 


doubt  as  to  the  existence  of  such  a  compound.  This  curve  has 
a  maximum  at  the  point  C2  (Fig.  58),  corresponding  to  20  atomic 
per  cent,  sulphur,  or  to  the  molecular  formula  SCl^  The  points 
EI,  E2,  E$  are  the  eutectic  points  for  S  +  S2C12,  S2C12  +  SC14  and 
SC14-1-C12,  respectively,  while  the  maximum  C\  corresponds  to 
the  compound  S2C12. 

Supersaturated  solutions. — A  sodium  sulphate  solution  satu- 
rated a  little  below  33°  can,  if  carefully  guarded  from  contact  with 
any  of  the  solid  salt,  be  cooled  down  to  room  temperature  without 
anything  crystallizing  out,  but  contact  with  the  tiniest  crystal 


346  INORGANIC  CHEMISTRY.  [§5  237 

fragment  of  Na2S04-10H2O  is  sufficient  to  cause  a  sudden  crystalli- 
zation of  this  salt. 

Sodium  sulphate  is  only  one  of  a  large  number  of  salts  capable 
of  forming  solutions  of  this  nature.  Sodium  thiosulphate  and  many 
of  the  nitrates  are  other  good  examples.  Such  solutions  are 
called  supersaturated.  They  are  perfectly  stable;  neither  rubbing 
with  a  glass  rod  nor  shaking  (which  treatment  ordinarily  tends 
to  induce  crystallization)  causes  the  formation  of  crystals,  provided 
no  trace  of  the  solid  salt  comes  in  contact  with  the  solution.  Such 
a  system,  which  is  unstable  under  only  one  condition,  is  called  a 
metastable  system.  (See  page  108.) 

If  a  supersaturated  solution  of  sodium  sulphate  is  cooled  down 
below  room  temperature,  another  hydrate  crystallizes  out,  viz., 
Na2SO4-7H20;  the  resulting  system  is  still  metastable,  however, 
for  contact  with  the  slightest  trace  of  Na2S04-10H2O  suffices  to 
convert  it  entirely  into  the  stable  system,  with  the  deposition  of 
Na2S04-10H20. 

The  smallest  amount  of  salt  (crystal  nucleus)  that  is  sufficient  to 
disturb  and  thus  cause  the  disappearance  of  a  metastable  system,  such 
as  is  represented  by  a  supersaturated  solution,  is  a  quantity  of  about 
the  order  10~10  g.>  according  to  OSTWALD.  The  extreme  minuteness 
of  this  amount  explains  why  a  spontaneous  disappearance  of  the  meta- 
stable condition  was  formerly  regarded  as  possible.  Inasmuch  as  very 
small  bits  of  crystals  are  always  floating  in  the  air  (especially  in  labora- 
tories), it  is  usually  only  necessary  to  open  a  bottle  containing  a  super- 
saturated solution  or  to  rub  the  sides  with  a  glass  rod  (which  always 
has  crystal  fragments  on  its  surface) ,  in  order  to  excite  crystallization 
into  the  stable  system. 

238.  For  the  reasons  stated  in  §§  65  and  66  it  is  assumed  that 
acids,  bases  and  salts  in  aqueous  solution  are  split  up  into  ions. 
This  dissociation  can  be  more  or  less  complete,  according  to  the 
nature  of  the  solute,  the  temperature  of  the  solution  and  its  con- 
centration. Examples  of  this  have  already  been  mentioned  here 
and  there  in  the  text;  hydrochloric  and  nitric  acids  in  tenth-normal 
solutions  are  almost  completely  dissociated,  carbonic  and  silicic 
acids  scarcely  at  all.  Among  the  bases  the  hydroxides  of  potas- 
sium, sodium  and  the  alkaline  earth  metals  are  almost  completely 
dissociated  at  this  dilution.  A  similar  difference  is  shown  by 


238.] 


SOLUTIONS. 


347 


salts ;  those  of  the  alkalies  are  practically  completely  ionized,  while 
mercuric  chloride  is  very  slightly  so.  We  shall  return  to  this  sub- 
ject more  in  detail  in  the  discussion  of  the  metals. 

The  principle  that  solutions  containing  equivalent  amounts  of  different 
electrolytes  differ  greatly  in  conductivity  and  hence  in  degree  of  ionization 
can  be  demonstrated  in  an  elegant  manner  with  the  aid  of  an  apparatus 
devised  by  WHITNEY  (Fig.  59). 


FIG.  59. 

Four  glass  cylinders  (3  cm.  diam.)  are  fitted  each  with  two  horizontal 
platinum  disks  (copper  can  be  used  but  is  less  satisfactory)  to  serve  as 
electrodes,  the  upper  ones  being  movable.  Each  lower  electrode  is 
connected  with  an  incandescent  lamp  and  the  apparatus  as  a  whole 
with  the  terminals  of  a  (preferably)  alternating  110-volt  circuit.  After 
placing  in  each  of  the  tubes  120  cc.  distilled  water  they  are  filled 
with  5  cc.  of  half-equivalent-normal  hydrochloric,  sulphuric,  monochlo- 
racetic  and  acetic  acids  respectively.  On  making  the  distance  between 
the  electrodes  alike  in  all  the  cylinders,  the  lamp  beneath  the  hydro- 
chloric acid  is  found  to  glow  brightest,  since  the  resistance  of  this  solu- 
tion is  the  least.  The  other  lamps  follow  in  brightness  in  the  order 
given  above.  The  electrodes  are  next  adjusted  so  that  all  the  lamps 
are  equally  bright,  when  it  is  seen  that  while  the  electrodes  in  the  hydro* 


348  INORGANIC  CHEMISTRY.  [§§  238 

chloric  acid  are  farthest  apart,  those  in  the  acetic  acid  are  almost  in 
contact. 

In  order  to  show  that  the  alkali  salts  of  these  acids,  unlike  the  acids 
themselves,  have  nearly  the  same  conductivities  and  degrees  of  disso- 
ciation the  solutions  are  just  neutralized  with  potassium  hydroxide 
and  the  lamp  test  repeated.  The  lamps  are  equally  brilliant  when 
the  electrodes  are  at  approximately  the  same  height. 

239.  In  the  solution  of  an  extensively  ionized  salt  we  should 
expect  to  find  the  properties  of  the  cation  and  the  anion.  It  must 
exhibit  the  sum  of  the  properties  of  the  two  ions,  or,  to  use  other 
words,  its  properties  must  be  additive  with  reference  to  those  of 
both  ions.  This  is  actually  the  case,  both  physically  and  chemically. 
As  for  the  chemical  properties,  we  observe  that  solutions  of  different 
salts  of  the  same  metal  all  give  the  same  reactions ;  from  the  solutions 
of  all  lead  salts,  for  instance,  hydrogen  sulphide  precipitates  black 
lead  sulphide,  sulphuric  acid  white  lead  sulphate,  etc.  Similarly 
the  solutions  of  salts  of  the  same  acid  are  all  characterized  by  the 
same  reactions;  sulphates,  for  example,  by  the  precipitate  they  give 
with  barium  chloride  solution.  All  this  appears  very  strange 
when  we  recall  that  the  solid  salts  are  markedly  different  from 
each  other  in  their  properties,  but  we  are  forced  to  just  such  a  con- 
clusion when  we  assume  that  the  salts  are  ionized  in  solution. 

Among  the  physical  properties  additivity  is  very  apparent  in 
the  color  of  salts  of  colored  acids  and  bases.  OSTWALD  found  that 
all  permanganates  with  a  colorless  base,  when  prepared  in  equiva- 
lent solutions,  give  exactly  the  same  absorption  spectrum  (§  263). 
All  dilute  copper  solutions  are  blue.  In  the  permanganates  it  is 
the  anion  MnO4',  in  the  copper  solutions  the  cation  Cu",  which 
is  to  be  regarded  as  the  color-carrier. 

When  the  solvent  is  one  in  which  ionization  does  not  occur,  the 
salts  of  the  same  base  may  differ  widely  in  color.  For  instance,  a  solu- 
tion of  cobalt  nitrate  in  alcohol  is  purple,  that  of  the  chloride  is  bluish 
violet;  but  if  both  are  poured  into  an  excess  of  water,  the  solution 
becomes  pink  in  each  case.  Another  example  is  found  in  the  alcoholic 
solutions  of  cupric  chloride  and  nitrate;  the  former  is  dark  green,  the 
latter  blue;  on  the  addition  of  water  both  become  blue. 

This  additive  nature  manifests  itself  in  various  other  physical 
properties  also.  But  since  it  cannot  usually  be  shown  directly 


239.J  SALT  SOLUTIONS.  349 

(as  in  the  case  of  colored  salts),  we  have  to  approach  the  matter 
somewhat  indirectly,  as  the  following  example  will  illustrate. 
The  specific  gravity  of  a  sugar  solution  can  be  represented  fairly 
accurately  by  the  formula 


n  being  the  number  of  moles  per  liter  and  K  a  constant.  Similarly 
in  the  case  of  the  solution  of  a  highly  ionized  salt  whose  specific 
gravity  is  raised  to  l+an  by  the  anion,  to  1  +pn  by  the  cation, 
a  and  /?  being  constants,  the  specific  gravity  of  the  solution,  if 
we  assume  additivity  to  exist,  must  be 


The  values  of  a  and  /?  are  as  yet  unknown.     For  salts  with  the 
same  anion  the  specific  gravity  is  expressed  by 


For  salts  with  the  same  anions  as  in  the  former  case  but  with  a 
different  cation  the  specific  gravities  of  their  solutions  are  repre- 
sented thus: 

Si'=l+w(ai+/?i);    S2'=l+n(ai+p2);    S3'=l  +n(>i  +&),   etc., 

whence  it  follows  that  the  differences  Si  —  Si',  82—82',  Sz  —  Ss'r 
=  n(a  —  ai),  must  always  have  the  same  value  in  case  additivity 
really  exists.  The  equality  of  these  differences  can  therefore  be 
used  as  a  proof  of  additivity. 

A  concrete  example  of  the  above  reasoning  is  to  be  found  in 
the  specific  gravity  values  of  the  solutions  KC1,  NaCl,  NH4C1  and 
KBr,  NaBr,  NH4Br.  Here  we  actually  have  the  relationship: 

KC1-  KBr  -  NaCl-  NaBr  =  NH4C1  -NH4Br  ; 

however,  not  simply  in  specific  gravity  but  with  reference  to  other 
physical  constants  as  well.  Compressibility,  capillarity  and  re- 
fractive index,  for  example,  have  been  found  to  conform  to  this 
same  additive  scheme. 

The  ionization  hypothesis  also  leads  us  to  predict  that  when 
dilute  solutions  of  strong  acids  and  bases,  each  containing  one 


350  ,       INORGANIC  CHEMISTRY.  [§§  239- 

mole,  are  mixed,  the  same  caloric  effect  will  be  observed.  This 
is  found  to  be  the  case  (13.8  cal.).  The  only  change  that  takes 
place  in  the  mixing  is  the  formation  of  water  from  its  ions  (§  66). 
Further,  the  so-called  law  of  thermo-neutrality,  which  says  that 
when  two  dilute  salt  solutions  are  mixed  there  is  no  caloric  effect, 
is  a  natural  consequence;  for  the  ions  of  the  two  salts  exist  in 
the  free  state  both  before  and  after  the  mixing. 

ACIDIMETRY  AND  ALKALIMETRY.    THEORY  OF 
INDICATORS. 

240.  The  amount  of  acid  or  base  present  in  a  liquid  can  be 
determined  most  simply  by  volumetric  analysis  (§  93).  Those  parts 
of  volumetric  analysis  which  comprise  the  methods  used  for  this 
purpose  are  known  as  acidimetry  and  alkalimetry.  Suppose  that 
we  wish  to  determine  the  amount  of  hydrochloric  acid  present  in  a 
given  volume  of  liquid.  A  known  volume  of  this  liquid  (50  cc., 
10  cc.,  or  less,  according  to  the  supposed  concentration)  is  meas- 
ured out  and  sodium  hydroxide  solution  of  known  concentration  is 
slowly  added  from  a  burette.  When  the  point  has  been  found  at 
which  the  liquid  becomes  neutral,  it  is  easy  to  calculate  the  con- 
centration of  the  acid  from  the  number  of  cubic  centimeters  of 
sodium  hydroxide  consumed. 

Example.  Determine  the  amount  of  nitric  acid  present  in  a  liter  of  a 
solution  of  this  acid  if  10  cc.  are  neutralized  by  7.3  cc.  of  a  normal  alkali 
solution.  These  7.3  cc.  are  equivalent  to  the  same  number  of  cubic  cen- 
timeters of  normal  nitric  acid.  Therefore  the  10  cc.  contain  7.3  milli- 
gram molecules  of  nitric  acid  or  63x7.3  mg.  One  liter  must  contain 
a  hundred  times  as  much,  or  45.99  g. 

Before  we  can  determine  the  concentration  of  an  acid  or  an 
alkali  in  this  manner,  we  must  first  possess  an  alkali  or  base  solu- 
tion of  known  concentration  and  further  have  a  delicate  means  of 
detecting  when  the  liquid  is  exactly  neutralized. 

1.  Preparation  of  an  acid  and  an  alkali  of  known  concentration. 
This  can  be  done  in  various  ways.  Oxalic  acid,  C2H2O4-2H2O, 
succinic  acid,  C4H6O4,  or  tartaric  acid,  C4H606?  can  be  used  as  the 
basis,  for  all  of  these  are  crystallized  solids  and  can  be  easily  obtained 
in  a  state  of  sufficient  purity;  hence  the  amount  of  acid  dissolved 


240,]  ACIDIMETRY  AND  ALKALIMETRY.  351 

can  be  very  accurately  determined  by  previously  weighing  the 
substance  on  an  analytical  balance.  We  thus  weigh  out  1  g.-equiv- 
alent  (J  g.-mol.)  of  one  of  these  acids,  dissolve  it  in  water  and 
dilute  to  exactly  a  liter.  Thereupon  with  the  help  of  this  normal 
acid  a  normal  alkali  is  prepared;  a  little  more  than  1  or  ^  or  ^, 
etc.,  gram-equivalent  of  sodium  hydroxide  or  potassium  hydroxide 
(barium  hydroxide  is  also  very  satisfactor}^)  is  dissolved  in  water 
and  this  solution  is  standardized  according  to  the  normal  acid,  i.e. 
the  concentration  is  determined  by  titration  with  normal  acid  and 
then  diluted  so  that  it  is  just  normal. 

Sodium  carbonate  can  also  be  used  as  a  basis.  After  being  first 
heated  in  order  to  expel  all  moisture  it  is  weighed  out  and  dissolved  in 
water.  This  solution  is  heated  to  boiling  and  covered  with  a  glass 
plate  with  a  hole  in  it,  through  which  the  nozzle  of  a  burette  is  passed. 
The  solution  of  the  acid  whose  concentration  is  to  be  determined  is 
then  allowed  to  flow  from  the  burette  into  the  boiling  liquid  till  neu- 
tralization is  effected.  Carbon  dioxide  escapes,  but  the  glass  plate 
prevents  any  loss  of  the  liquid  by  spurting. 

The  standardizing  can  also  be  accomplished  by  adding  the  acid 
solution  that  is  to  be  standardized  to  a  mixed  solution  of  potassium 
iodide  and  potassium  iodate.  Hydriodic  and  iodic  acids  are  set  free 
and  they  react  at  once  in  the  following  manner: 

5KI +  KI03  +  6HX  =5HI  +HI03  +  6KX;   SHI  +HI03  =3H20  +  61. 

Thus  for  every  equivalent  of  acid  one  atom  of  iodine  is  set  free.  By 
titrating  with  sodium  thiosulphate  the  amount  of  iodine  liberated  can 
be  determined.  This  method  gives  very  accurate  results. 

2.  Determination  of  the  point  at  which  the  liquid  becomes  neutral. 
Since  the  point  of  neutralization  of  an  acid  by  a  base  or  vice  versa, 
is  not  indicated  by  any  visible  phenomena,  a  minute  quantity  of 
some  substance  is  added  whose  color  is  altered  by  an  excess  of  the 
neutralizing  liquid.  Such  substances  are  litmus  (blue  in  alkaline 
and  red  in  acid  solutions),  phenolphthalem  (red  in  alkaline,  color- 
less in  acid  solutions),  methyl  orange  (yellow  in  alkaline,  red  in  acid 
solutions),  and  many  others.  Therefore,  on  gradually  adding  an 

a    a  11ne  solution  to  an  ^^. —  solution  in  the  presence  of  one  of 
acid  alkaline 


352  INORGANIC  CHEMISTRY.  [§§240^ 

these  substances  a  change  of  color  will  be  noticed  when  the  point  of 
neutrality  is  just  passed.  Coloring-matters  like  the  above  are 
termed  indicators.  The  change  of  color  is  due  in  many  cases  to  a 
transformation  of  the  substance  into  a  salt  whose  free  acid  is  very 
unstable  and  passes  over  almost  immediately  into  an  isomer  hav^ 
ing  a  different  color  from  the  free  acid  or  the  salt. 

241.  From  the  standpoint  of  the  ionic  theory  the  following 
theory  of  indicators  has  been  advanced:  If  a  couple  of  drops  of 
the  indicator  are  introduced  into  an  acid  solution,  the  ionization 
of  the  indicator,  which  is  only  very  slight,  is  reduced  by  the  great 
excess  of  acid  to  practically  zero.  If  a  base  is  then  added,  the 
H-ions  of  the  acid  to  be  titrated  are  removed  by  the  OH-ions. 
However,  if  the  acid  is  very  strong,  enough  H-ions  remain  in  the 
liquid  up  to  the  last  to  prevent  anything  like  an  extensive  ioniza- 
tion of  the  coloring-substance;  not  until  the  first  excessive  drop  of 
alkali  is  added  do  the  anions  of  the  coloring-substance  come  into 
existence,  the  alkali  compound  of  the  latter  being  strongly  disso- 
ciated. The  change  of  color  is  therefore  sharply  defined,  for  it  is 
due  to  this  difference  in  color  of  the  non-ionized  molecule  and  the 
anion.  On  the  other  hand,  if  the  acid  is  a  weak  one,  there  will  not 
be  enough  H-ions  present  when  the  end  of  the  titration  is  nearly 
reached  to  prevent  a  slight  ionization  of  the  coloring-substance. 
As  a  result  we  shall  have  in  the  solution  not  only  the  undissociated 
coloring-substance  but  its  anions  as  well,  even  before  the  titration 
is  completed, — in  other  words,  the  change  of  color  becomes  more 
gradual  and  hence  the  end  reaction  more  indefinite.  The  effect 
will  be  the  same  if  the  alkali  employed  contains  carbonate.  In 
that  case  near  the  end  of  the  titration  the  solution  will  only  contain 
carbonic  acid,  which  is  very  weak;  consequently  the  color  change 
will  not  be  sudden.  It  is  for  this  reason  that  in  titrating  soda  solu- 
tions (see  §  240)  the  carbonic  acid  must  be  expelled  by  boiling. 

If  a  weak  acid  is  to  be  titrated,  it  is  necessary,  according  to  the 
above,  to  select  an  indicator  which  is  much  less  ionized  even  than 
the  acid  itself  and  whose  alkali  salts  are  sufficiently  ionized  to 
produce  a  distinct  change  of  color.  A  very  suitable  one  for  this 
purpose  is  phenolphthalem.  Acetic  acid,  for  example,  can  be  sat- 
isfactorily titrated  with  it,  if  a  strong  base  is  employed,  for  the 
reasons  set  forth  above.  On  the  other  hand,  in  case  a  weak  base 
is  to  be  titrated,  phenolphthalein  is  not  so  satisfactory.  Ammonia 


242.]  COPPER.  353 

does  not  color  a  phenolphthale'in  solution  till  a  considerable  excess 
is  added,  because  at  the  great  dilution  in  which  the  ammonium- 
phenolphthalein  compound  exists  in  a  titration  it  is  almost  com- 
pletely split  up  by  hydrolysis  (§66). 

If  a  weak  base  is  to  be  titrated,  an  indicator  must  be  selected 
which  is  a  relatively  strong  acid,  for  then  the  salt  of  the  coloring- 
substance  will  be  hydrolyzed  only  to  a  limited  extent,  even  near 
the  termination  of  the  titration  (i.e.  when  the  concentration  of 
the  base  has  become  weak),  and  the  color  of  its  ions  will  therefore 
still  predominate.  For  such  a  titration  a  strong  acid  (e.g.  hydro- 
chloric or  sulphuric  acid)  must  be  used,  in  order  that  the  first  drop 
after  the  point  of  neutralization  is  reached  may  diminish  the  elec- 
trolytic dissociation  of  the  coloring-substance  and  so  give  the  solu- 
tion the  color  of  the  non-ionized  molecules.  Methyl  orange  is  an 
indicator  that  answers  these  requirements;  it  serves  very  well  in 
the  titration  of  ammonia.  All  other  indicators  are  intermediate  to 
these  two  extremes  (phenolphthalem  and  methyl  orange)  as  re- 
gards their  ionization,  and  their  applicability  is  determined  accord- 
ingly. 

COPPER. 

242.  This  metal  occurs  native  in  America,  China  and  Japan, 
forming  regular  crystals.  Other  copper  minerals  are  cuprite 
(Cu20),  malachite  and  azurite  (both  basic  carbonates),  chalcocite 
(Cu2S)  and  particularly  chalcopyrite,  or  copper  pyrites  (CuFeS2). 
The  United  States  furnishes  about  60%  of  the  world's  copper 
supply. 

The  extraction  of  the  metal  from  non-sulphurous  ores  is 
very  simple.  They  are  smelted  with  coal  and  thus  reduced  to  the 
metallic  state.  If  the  copper  ore  contains  sulphur,  the  metal- 
lurgical process  is  much  more  complicated  and  has  numerous  modi- 
fications. The  ore  is  broken  up  and  "  calcined  "  so  as  to  convert 
some  of  the  copper  sulphide  into  copper  oxide.  Thereupon  it  is 
fused  with  sand  and  other  siliceous  fluxes  (as  well  as  coal  for 
reducing  copper  sulphate),  and  the  iron,  but  not  the  copper,  is 
converted  into  silicate.  The  object  of  the  flux,  here  as  with  other 
metals,  is  to  lower  the  fusing  temperature  of  the  ore  and  collect  the 
impurities  (iron  in  this  case)  into  a  "  slag  "  consisting  of  fused 


354  INORGANIC  CHEMISTRY, 

silicates,  etc.  The  slag  floats  and  can  be  run  off.  The  fusion 
process  is  repeated  until  all  the  iron  is  eliminated.  The  resulting 
mixture  of  impure  copper  sulphide  and  copper  oxide  is  called  matte 
(also  regulus  and  coarse  metal).  By  repeated  roasting  and  fusing, 
crude  metallic  copper  is  obtained: 

2Cu2O  +  Cu2S  =  6Cu + S02. 

Finally  it  is  fused  with  coal  to  reduce  any  copper  oxide  remaining. 

Refining.  The  copper  thus  obtained  often  contains  small  quantities 
of  other  metals.  Since  these  impurities  lower  its  conductivity,  a  better 
grade  is  demanded  for  electrical  purposes.  Crude  copper  is  now  refined 
by  an  electrolytic  process  which  yields  chemically  pure  copper.  If  an 
impure  copper  solution  is  electrolyzed,  it  is  possible  under  suitable  con- 
ditions to  precipitate  pure  copper  in  a  compact  mass,  while  the  impuri- 
ties remain  in  solution  or  are  deposited  as  powder.  From  this  powder 
("slimes  ")  a  considerable  amount  of  gold  and  silver  is  obtained. 

The  usual  arrangement  is  to  suspend  plates  ("anodes")  of  crude 
copper  and  thin  sheets  of  pure  copper  alternately  in  a  copper  vitriol 
solution  acidified  with  sulphuric  acid.  If  the  crude  plates  are  then 
connected  with  the  positive  pole  and  the  thin  sheets  with  the  negative 
pole  of  the  dynamo  current,  pure  copper  is  deposited  on  the  sheets, 
while  an  equivalent  amount  of  the  crude  copper  dissolves  to  take  its 
place. 

Physical  Properties.  Copper  has  a  bright  red  color.  It  is 
rather  hard  but  very  extensible  and  flexible;  it  can  be  drawn  out 
into  very  fine  wire  and  beaten  into  extremely  thin  sheets  (imitation 
gold-leaf),  which  are  green  in  transmitted  light.  Sp.  g.  =  8.94; 
melting  point  =1083°;  boiling-point,  2310°. 

Chemical  Properties.  In  dry  air  copper  is  permanent  at  ordi- 
nary temperatures,  but  in  moist  air  it  becomes  covered  with  a  thin 
coating  of  basic  copper  carbonate,  which  protects  it  from  further 
rusting.  On  being  heated  in  the  air  it  turns  to  copper  oxide,  CuO. 
It  is  readily  attacked  by  nitric  acid  (§  120),  but  not  by  dilute  hydro- 
chloric acid.  Sulphuric  acid  has  no  effect  on  it  at  ordinary  tem- 
peratures, but  at  higher  temperatures  a  reaction  takes  place  in 
which  sulphur  dioxide  is  given  off  (§  78).  Ammonia  and  oxygen 
dissolve  it  to  form  a  blue  liquid/  copper  oxide  ammonia.  Copper 
is  deposited  from  solutions  of  its  salts  by  iron,  magnesium  and 
other  metals. 


243.]  COMPOUNDS  OF  COPPER.  355 

Uses  and  alloys.  Copper  finds  extensive  use  in  the  arts,  both 
as  such  and  in  alloys.  The  well-known  yellow  b  r  a  s  s  is  an  alloy 
of  1  part  zinc  and  2  parts  copper  and  is  harder  than  copper  itself. 
German  silver  consists  of  about  50%  Cu,  25%  Ni,  and 
25%  Zn;  its  electrical  conductivity  is  affected  very  little  by  changes 
of  temperature,  which  makes  it  valuable  for  resistance  coils,  etc. 
For  bronzes  see  §  199. 

Copper  is  employed  in  large  quantities  inelectrotyping. 
A  cast  is  first  constructed  of  plaster  of  Paris  and  made  a  conductor 
by  being  coated  with  graphite,  whereupon  it  is  suspended  by  the 
wire  of  a  battery  into  a  copper  sulphate  solution;  a  plate  of  pure 
copper  serves  as  the  anode.  If  the  potential  difference  at  the 
electrodes  is  properly  regulated,  the  copper  is  deposited  on  the 
plaster  cast  in  compact  form,  so  that  all  the  details  of  the  original 
are  reproduced  with  the  greatest  fidelity. 

Compounds  of  Copper. 

243.  Copper  forms  two  sets  of  salts,  which  are  derived  from  the 
oxides  Cu2O,  cuprous  oxide,  and  CuO,  cupric  oxide. 

CUPROUS   COMPOUNDS. 

Cuprous  oxide,  Cu2O,  can  be  obtained  from  cupric  salts  in 
various  ways,  e.g.  by  reducing  them  in  alkaline  solution  with  grape 
sugar,  hydroxylamine,  arsenious  acid,  or  the  like.  It  forms  a 
reddish-yellow  crystalline  powder,  which  is  unaffected  by  the  air 
at  ordinary  temperatures.  When  cuprous  oxide  is  heated,  it 
breaks  up  into  cuprous  oxide  and  oxygen;  2CuO^±Cu2O  +  O. 
At  1025°  the  dissociation  tension  of  the  cupric  oxide  reaches 
150  mm. ;  consequently  at  this  temperature  in  the  air  it  passes 
over  completely  into  cuprous  oxide,  since  the  partial  pressure 

of  the  oxygen  of  the  air  is =  152  mm.     Cuprous  oxide  dis- 

5 

solves  in  ammonia;  this  solution  rapidly  turns  blue  because  of 
the  absorption  of  oxygen,  the  cuprous  oxide  going  over  into 
cupric  oxide.  Cuprous  oxide  is  transformed  by  sulphuric  acid 
into  copper  sulphate  and  copper: 

Cu20  +  H2SO4  =  CuSO4 + Cu  +  H20. 


356  INORGANIC    CHEMISTRY.  [§  243. 

It  is  possible  that  cuprous  sulphate  is  first  formed 
and  that  the  cuprous  ions  of  this  solution  are  forthwith 
changed  into  cupric  ions  and  non-ionized  copper: 


Of  the  cuprous  salts  the  sulphate  and  halides  are  known. 
Cu2Cl2,  Cu2Br2,  and  Cu2I2  are  all  "  insoluble  "  (cf.  §  235)  ;  their 
solubility  decreases  with  increasing  atomic  weight  of  the 
halogen. 

Cuprous  chloride,  Cu2Cl2  (the  vapor  density  indicates  this 
doubled  formula),  separates  out  when  a  solution  of  cupric  chloride 
is  boiled  with  copper,  or  when  a  mixed  solution  of  copper  sulphate 
and  sodium  chloride  is  saturated  with  sulphur  dioxide  gas  and  the 
resulting  liquid  poured  into  water.  It  is  a  white  crystalline  sub- 
stance, which  must  be  kept  under  water,  for  it  absorbs  oxygen 
rapidly  when  moist  and  turns  green  because  of  the  formation  of 
basic  copper  chloride,  CuCl-OH.  It  melts  at  430°  and  distils  at 
about  1000°.  It  is  soluble  in  concentrated  hydrochloric  acid  and 
in  ammonia.  These  solutions  are  at  first  colorless  but  very  soon 
become  blue  because  of  the  absorption  of  oxygen  (formation  of 
cupric  compounds).  They  also  have  the  power  of  absorbing  car- 
bon monoxide,  forming  an  unstable  compound,  Cu2Cl2  •  CO  •  2H2O 
which  crystallizes  in  colorless  laminae.  Use  is  made  of  this  prop- 
erty in  gas  analysis. 

Cuprous  iodide,  Cu2I2,  is  formed  when  a  solution  of  copper 
sulphate  is  treated  with  potassium  iodide,  half  of  the  iodine  being 
liberated: 

2CuS04  +4KI  =2K2S04  +Cu2I2  +I2. 

It  may  be  supposed  that  cupric  iodide  is  first  formed  and  that 
it  then  breaks  up  into  cuprous  iodide  and  iodine,  or,  rather,  that  the 
ions  of  cupric  iodide  interact  thus: 

Cu"+2I'=CuI+I, 
the  cuprous  iodide  being  unionized  because  "insoluble." 

According  to  OSTWALD,  however,  an  equilibrium  is  formed  here, 
for,  though  the  cuprous  iodide  is  but  slightly  soluble,  the  reaction  does 
not  complete  itself  and  some  cupric  ions  still  remain  in  solution.  The 


§243.]  CUPROUS  COMPOUNDS.  357 

reversibility  of  this  reaction  is  evident  from  the  fact  that  cuprous  iodide 
is  dissolved  by  an  alcoholic  iodine  solution,  so  that  we  have 

Cu"  +  2I'<=±CuI+I. 

Therefore,  in  order  to  make  the  precipitation  more  complete,  a  sub- 
stance (S02)  is  added,  which  will  remove  the  iodine,  one  of  the  reaction 
products.  This  treatment  is  especially  effective  because  the  iodine  is 
thereby  converted  into  ions  and  this  raises  the  concentration  of  one 
of  the  components  on  the  left  side  of  the  equilibrium  equation. 

Cuprous  cyanide,  Cu2(CN)2,  can  be  obtained  in  a  manner  analo- 
gous to  that  described  for  cuprous  iodide,  viz.,  by  mixing  solutions 
of  copper  sulphate  and  potassium  cyanide.  Half  of  the  cyanoge^ 
escapes  as  gas: 

2CuS04  +4KCN=2K2S04  +Cu2(CN)2  +  (CN)2. 

Cuprous  cyanide  dissolves  very  rapidly  in  an  excess  of  potas- 
sium cyanide,  forming  a  salt,  2KCN-Cu2(CN)2,  which  contains  a 
complex  anion  [Cu2(CN)4]".  Practically  all  of  the  copper  ions  go 
to  form  these  complex  ions  on  the  addition  of  potassium  cyanide, 
for  the  solution  gives  none  of  the  ordinary  reactions  for  copper, 
not  even  that  with  hydrogen  sulphide,  although  copper  sulphide 
is  precipitated  by  this  reagent  even  when  the  concentration  of  the 
copper  ions  is  very  slight  (§  73). 

Cuprous  sulphate,  Cu2S04,  is  formed  by  the  action  of  methyl 
or  ethyl  sulphate  on  cuprous  oxide  in  the  absence  of  water  at 
160°: 

Cu20  +  Me2SO4=  Cu2SO4+ Me2O. 

Water  decomposes  it  rapidly  with  evolution  of  heat : 

Cu2SO4  (solid)  +  Aqua  =  CuS04  (dissolved)  +  Cu  (solid) +21  cals. 

f 

This  explains  (§  101)  why  many  attempts  to  prepare  the  salt 
resulted  in  failure  and  it  was  for  a  long  time  thought  incapable 
of  preparation. 


358  INORGANIC  CHEMISTRY.  [§§244- 


CUPRIC   COMPOUNDS. 

244.  Cupric  oxide,  CuO,  is  a  dense,  black  powder,  obtained  by 
heating  copper  in  the  presence  of  oxygen  at  a  high  temperature.  It 
can  also  be  prepared  by  heating  the  nitrate  to  redness  or  igniting 
the  hydroxide  or  the  carbonate.  When  finely  divided  it  occludes 
on  its  surface  large  amounts  of  steam.  It  finds  extensive  use  in 
organic  analysis. 

^Cupric  hydroxide,  CuO-nH20,  separates  out  as  a  flocculent, 
voluminous  blue  precipitate  (hydrogel,  §  195)  when  the  solution 
of  a  copper  salt  is  treated  with  caustic  potash  or  soda.  On  boiling 
the  liquid  it  turns  black,  water  being  liberated  and  cupric  oxide 
formed. 

Cupric  chloride,  CuCl2  •  2H20,  is  obtained  by  dissolving  cupric 
oxide  or  carbonate  in  hydrochloric  acid.  It  crystallizes  in  blue 
rhombic  needles,  which,  however,  often  appear  green  because  of 
mother-liquor  adhering  to  them.  It  is  readily  soluble  in  water 
and  alcohol.  The  anhydrous  salt  is  yellow;  the  concentrated 
aqueous  solution  is  green;  the  dilute  solution  is  blue.  This  differ- 
ence can  be  attributed  to  the  breaking  up  of  the  salt  into  its  ions, 
for  all  dilute  copper  solutions  are  blue,  no  matter  what  the  acid 
radical  is.  It  therefore  follows  that  the  copper  ion  imparts  a  blue 
color  to  solutions.  The  green  color  of  a  concentrated  solution 
seems  most  probably  due  to  the  formation  of  complex  ions, 
such  as  (CuCl3)'. 

Cupric  bromide  is  analogous  to  the  chloride;  cupric  iodide  is 
unstable,  decomposing  at  once  into  iodine  and  cuprous  iodide 
(§  243). 

Copper  sulphate,  CuS04  •  5H20,  blue  vitriol,  is  the  most  familiar 
salt  of  copper.  It  is  obtained  as  a  by-product,  chiefly  from  gold 
and  silver  refineries,  and  is  also  manufactured  by  dissolving 
copper  in  sulphuric  acid.  It  crystallizes  in  large  blue  triclinic 
crystals,  which  lose  four  molecules  of  water  at  100°;  the  fifth 
is  liberated  at  200°.  The  anhydrous  copper  sulphate  is  a  white 
powder,  which  absorbs  water  greedily,  turning  blue  again.  At 
20°  100  parts  H20  dissolve  42.31  parts  of  the  crystallized  sulphate. 
Blue  vitriol  is  employed  in  large  quantities  in  electroplating,  etc. 
(§  242). 


245-]  SILVER.  359 

Copper  nitrate,  Cu(N03)2,  can  crystallize  with  three  or  six 
molecules  of  water  and  is  dark  blue. 

Copper  carbonate.  The  normal  salt  is  unknown,  but  basic  salts 
have  been  prepared. 

Copper  arsenite,  CuHAs03,  is  used  as  a  pigment  under  the 
name  of  Scheele's  green.  Schweinfurth  green,  or  Paris  green,  is  a 
double  compound  of  copper  arsenite  and  copper  acetate.  Since 
both  are  very  poisonous,  their  use  in  dyeing  textile  fabrics,  wall- 
paper, etc.  (§  157)  is  being  restricted. 

Copper  sulphide,  CuS,  is  formed  as  a  black  precipitate  by  pass- 
ing hydrogen  sulphide  into  a  copper  solution.  When  moist  it 
oxidizes  slowly  in  the  air  to  copper  sulphate.  On  being  heated 
in  a  current  of  hydrogen  it  yields  cuprous  sulphide,  Cu2S,  and 
hydrogen  sulphide. 

Copper  salts  and  ammonia.  On  mixing  a  solution  of  ammonia 
with  a  copper  salt,  a  precipitate  of  copper  hydroxide  is  first  formed, 
if  not  too  much  ammonia  is  used;  this  precipitate  is  dissolved  by 
an  excess  of  ammonia  to  a  dark  blue  solution.  If  the  latter  is 
evaporated  or  treated  with  alcohol,  ammoniacal  compounds  crys- 
tallize out;  a  typical  one  is  CuSO4-4NH3-H2O,  which  is  trans- 
formed into  CuSO4-2NH3  on  being  heated  to  150°.  The  aqueous 
solutions  of  these  substances  are  to  be  assumed  to  contain  complex 
ions  of  copper  and  ammonia,  since  they  do  not  give  some  of  the 
ordinary  copper  reactions,  e.g.  precipitation  with  caustic  potash. 
The  fact  that  certain  other  reactions  of  copper  do  however  appear, 
e.g.  precipitation  with  hydrogen  sulphide,  proves  that  free  copper 
ions  are  still  present  in  the  liquid,  although  only  to  a  small  extent. 

SILVER. 

245.  This  metal  occurs  native;  nuggets  weighing  100  kilos  are 
not  unknown.  The  important  silver  ores  are  argentite,  Ag2S, 
stromeyerite,  Cu2S-Ag2S,  pyrargyrite,  3Ag2S-Sb2S3,  and  stephanite, 
Ag5S4Sb.  It  is  also  found  in  smaller  amounts  in  cerargyrite,  or 
horn  silver,  AgCl.  Traces  of  silver  compounds  are  known  to  exist 
in  sea-water.  Many  lead  ores,  e.g.  galenite,  contain  a  small  per- 
centage of  silver  and  in  some  cases  it  is  extracted. 

The  chief  silver-producing  countries  are  the  United  States 
(Colorado  and  neighboring  States),  Mexico,  Australia  and  Bolivia. 


360  INORGANIC   CHEMISTRY.  [§  245. 

The  present  annual  world's  production  of  silver  is  about  55,000,000 
troy  ounces  (1,710,000  kg.). 

Silver  is  now  generally  obtained  from  its  ores  by  the  cyanide 
process.  In  this  process  the  pulverized  ore  is  allowed  to  stand 
for  some  time  in  a  weak  solution  (0.1-0.4%)  of  sodium  cyanide. 
The  silver  sulphide  dissolves  as  the  double  cyanide  according 
to  the  equation: 

Ag2S  +  4NaCN= 2AgNa(CN)2  +  Na2S. 

As  soon  as  a  certain  amount  of  silver  has  gone  into  solution,  an 
equilibrium  is  formed,  because  the  Na2S  tends  to  react  back- 
ward with  the  formation  of  silver  sulphide.  This  is  avoided 
by  blowing  air  into  the  solution,  and  thereby  oxidizing  the  sodium 
sulphide.  Metallic  silver  is  also  taken  up  by  the  sodium  cyanide 
solution  (§  248) ;  so  is  the  chloride,  or  horn  silver.  The  recovery 
of  the  silver  from  the  double  cyanide  solution  is  accomplished  by 
precipitation  with  zinc,  or  by  electrolytic  deposition. 

Lead  ores  usually  contain  some  silver.  In  the  smelting  of  lead 
the  silver  all  goes  into  the  lead  and  is  recovered  from  it  in  the 
following  way:  The  argentiferous  lead  is  fused  and  then  cooled 
slowly  till  it  begins  to  congeal.  Just  as  pure  ice  crystallizes  out  of 
a  dilute  salt  solution  on  cooling,  so  the  lead  separates  out  here 
in  crystals  free  from  silver.  These  are  removed  and  this  process — 
called  Pattinsonizing  after  its  inventor — is  kept  up  till  the  percent- 
age of  silver  reaches  about  1%.  This  rich  lead  is  then  subjected 
to  the  cupellation  process,  i.e.  the  lead  is  fused  in  a  rever- 
beratory  furnace,  whose  hearth  consists  of  a  porous  mass  (cupel, 
or  test).  The  lead  is  oxidized  to  the  easily  fusible  oxide  PbO 
(litharge)  which  is  partly  driven  off  by  a  blast  from  time  to  time 
through  the  channel  provided  for  its  escape,  and  partly  absorbed 
by  the  porous  material  (bone-ash,  or  clay  and  limestone)  of  the 
cupel.  Towards  the  end  of  the  process  the  film  of  litharge  remain- 
ing becomes  so  thin  that  the  silver  beneath  reflects  the  light,  pro- 
ducing a  beautiful  iridescence.  Here  and  there  the  film  breaks, 
disclosing  the  brilliant  surface  of  the  metal  ("  brightening  "  of  the 
silver).  The  silver  is  finally  left  in  the  metallic  state. 

Another  method  (PARKES')  involves  the  principle  of  distri- 
bution between  two  slightly  miscible  solvents  (ORG.  CHEM., 
§  24).  Molten  zinc  and  lead  are  only  slightly  soluble  in  each 
other.  Silver  is  several  hundred  times  more  soluble  in  molten 


§  245.]  SILVER.  361 

zinc  than  in  molten  lead;  thus  the  silver  can  be  very  fully 
extracted  from  the  lead  by  fusing  with  zinc.  The  process  is  as 
follows:  To  the  fused  argentiferous  lead  some  zinc  (containing 
0.5%  Al)  is  added  and  the  mixture  is  stirred.  The  zinc  takes 
up  most  of  the  silver  from  the  lead  and  floats  on  the  molten  mass. 
It  is  skimmed  off  and  cast  into  plates,  which  serve  as  the  anodes 
in  the  subsequent  electrolysis.  At  the  cathode  nearly  pure  zinc 
is  deposited,  while  silver  powder  (70-80%  Ag,  the  rest  Pb)  sinks 
to  the  bottom  and  is  removed  to  the  cupel. 

The  electrolytic  refining  of  silver  is  now  carried  on  exten- 
sively in  America.  A  great  deal  is  recovered  from  the  copper 
slimes  (§  242). 

The  pure  silver  of  commerce  usually  contains  a  little  copper 
and  other  metals;  STAS  obtained  it  chemically  pure  by  dissolving 
the  product  of  the  smelter  in  nitric  acid  and  precipitating  it  with 
hydrochloric  acid  as  the  chloride;  this  was  then  reduced  by  boiling 
with  dilute  caustic  potash  and  milk  sugar  and  finally  distilled  with 
the  aid  of  an  oxy hydrogen  flame  in  an  apparatus  made  of  lime. 

Physical  Properties.  Silver  crystallizes  in  regular  octahe- 
drons. It  has  a  white  color  and  a  high  lustre.  It  is  the  best 
conductor  of  heat  and  electricity  of  all  the  metals  and  it  is  very 
malleable  and  ductile.  Sp.  g.  =  10.5;  m.-pt.  =  961°;  boiling- 
point  =1955°.  It  volatilizes  in  the  form  of  a  blue  vapor  (STAS). 
Molten  silver  absorbs  oxygen,  but  allows  it  to  escape  on  becoming 
solid  (§  9). 

Chemical  Properties.  Silver  is  one  of  the  precious  metals; 
this  term  is  applied  chemically  to  those  metals  which  do  not  com- 
bine with  oxygen  directly  (under  ordinary  pressure)  either  at 
ordinary  or  higher  temperatures.  If,  however,  the  pressure  is 
raised,  silver  combines  with  oxygen  directly  at  an  elevated  tem- 
perature: 

2Ag+O±Ag20. 

Nitric  acid  attacks  it  readily  at  ordinary  temperatures,  sulphuric 
acid  only  at  higher  temperatures;  hydrochloric  acid  has  very  little 
effect  on  it. 

Uses;  alloys.  Pure  silver  is  seldom  made  use  of  practically, 
but  its  alloys  are  employed  in  the  manufacture  of  silverware  and 
coins.  For  these  purposes  an  alloy  with  copper  is  used.  Silvei 


362  INORGANIC  CHEMISTRY.  [§§245- 

plate  and  jewelry  usually  contain  75  or  more  per  cent  of  silver;  the 
silver  coins  of  the  United  States  and  continental  countries  consist 
of  90%  silver  and  10%  copper;  the  English  shillings  ("sterling" 
silver)  contain  92.5%  silver.  The  admixture  of  copper  makes  the 
metal  harder. 

Considerable  silver  is  consumed  in  silvering  objects  of  copper 
or  other  metals  (silver-plating).  At  present  this  is  usually 
done  by  electrolysis  (§  242).  The  object  to  be  plated  is  made  the 
cathode  and  a  silver  plate  the  anode;  the  bath  consists  of  silver 
cyanide  dissolved  in  an  excess  of  potassium  cyanide. 

The  market  price  of  silver,  on  a  gold  basis,  was  almost  constant  from 
1650  to  1870,  the  market  ratio  of  the  two  metals  at  London  remaining 
very  close  to  1:15.5,  but  in  the  last  few  decades  the  value  of  silver  has 
decreased  relatively  very  greatly.  The  reason  for  this  shrinkage  of 
value  is  to  be  found  not  only  in  the  chemical  improvements  in  the 
metallurgy  of  silver,  but  much  rather  in  the  discovery  of  large  deposits 
of  high-percentage  silver  ores,  especially  in  the  Western  Hemisphere. 
In  1902  the  ratio  was  1:39.15;  since  then  it  has  improved  somewhat, 
being  in  1911  1:38.74. 

Compounds  of  Silver. 

246.  The  known  oxides  are:  Ag40,  silver  suboxide  (very  un- 
stable); Ag20,  silver  oxide,  from  which  the  salts  of  silver  can  be 
derived;  and  AgO,  silver  peroxide  (formed  from  silver  and  ozone). 

Silver  oxide,  Ag20;  is  deposited  as  a  dark  brown  amorphous 
precipitate  when  the  solution  of  a  silver  salt  is  treated  with  caustic 
soda  or  baryta-water  free  from  carbonic  acid.  It  is  somewhat  sol- 
uble in  water  (2.16X10~4  mole  per  liter  at  25°) ;  the  solution  prob- 
ably contains  the  silver  hydroxide,  for  it  reacts  alkaline  and  must 
therefore  contain  hydroxyl  ions.  In  its  saturated  aqueous  solu- 
tion 70%  of  the  molecules  are  found  to  be  ionized.  Silver  hydrox- 
ide is  thus  not  so  strong  a  base  as  the  alkalies,  but  considerably 
stronger  than  ammonia.  Moist  silver  oxide  (AgOH)  absorbs  car- 
bon dioxide  from  the  air  and  the  silver  salts  react  neutral,  while  the 
salts  of  most  of  the  other  heavy  metals  give  an  acid  reaction  be- 
cause of  a  slight  hydrolytic  dissociation  in  aqueous  solution.  By 
heating  to  250°  silver  oxide  is  broken  up  into  its  elements.,  It  is 
reduced  by  hydrogen  at  as  low  a  temperature  at  100°.  Ammonia 


246-]  COMPOUNDS  OF  SILVER.  363 

water  dissolves  it  readily  because  of  the  formation  of  a  complex 
ion,  Ag(NH3)2* . 

Silver  chloride,  AgCl,  is  obtained  by  precipitating  a  silver  solu- 
tion with  hydrochloric  acid  or  a  soluble  chloride  like  sodium  chlo- 
ride; it  forms  a  characteristic  "  curdy"  precipitate.  It  is  almost 
insoluble  in  water,  1-  part  in  715,800  H20  at  13.8°. 

When  a  silver  solution  is  added  very  carefully  to  a  sodium  chloride 
solution  (or  to  another  chloride),  a  point  can  be  found  when  the  liquid 
gives  a  cloudiness  (due  to  AgCl)  with  either  solution.  This  must  be 
attributed  to  the  fact  that  the  liquid  is  saturated  with  silver  chloride  and 
contains  no  other  silver  salt  nor  any  other  chloride.  In  view  of  the  very 
high  dilution  of  such  a  silver  chloride  solution  (see  above)  it  may  be 
assumed  that  the  dissolved  part  is  completely  ionized.  If  silver  or 
chlorine  ions  are  now  introduced  into  the  liquid,  the  ionization  of  the 
silver  chloride  is  diminished  and  AgCl  molecules  are  formed,  but  these 
cannot  remain  in  solution,  since  the  solution  is  already  saturated  with 
them. 

Silver  chloride  dissolves  readily  in  ammonia,  potassium  cyanide 
and  sodium  thiosulphate,  forming  complex  ions. 

If  a  solution  of  silver  chloride  and  ammonia  is  allowed  to  evap- 
orate in  the  dark  at  room  temperature,  silver  chloride  crystallizes 
out  in  finely  developed  octahedrons. 

Silver  bromide,  AgBr,  is  less  soluble  than  the  chloride  and  has 
a  yellowish  color.  It  dissolves  with  difficulty  in  ammonia  but 
easily  in  thiosulphate.  Silver  iodide,  Agl,  is  even  less  soluble  than 
silver  bromide  at  ordinary  temperatures.  It  is  insoluble  in  ammo- 
nia. It  is  yellow.  At  high  temperatures  these  halides  melt  and 
on  cooling  form  a  horny  mass,  which  can  be  cut  with  the  knife 
("  horn-silver, "  cf.  §  245).  Silver  fluoride,  AgF,  is  much  more  solu- 
ble in  water  than  the  three  preceding  halogen  compounds. 

Potassium  silver  cyanide,  KAg(CN)2,  obtained  on  adding  potas- 
sium cyanide  to  a  silver  solution,  dissolves  readily  in  water  and  is 
used  in  large  quantities  in  electro-plating.  When  a  current  passes 
through  it,  potassium  is  deposited  (primarily)  at  the  cathode,  while 
the  anion  Ag(CNV  wanders  to  the  anode;  however,  potassium  pre- 
cipitates silver  from  potassium  silver  cyanide: 

K+KAg(CN)2=2KCN+Ag. 


364  INORGANIC  CHEMISTRY.  [§§  246- 

Thus  silver  is  deposited  on  the  cathode  while  the  anion  Ag(CN}/ 
takes  up  an  atom  of  silver  at  the  silver  anode  to  form  silver  cyanide 
and  again  unites  with  potassium  cyanide  to  form  the  double  salt; 
if  the  anode  is  of  platinum,  cyanogen  gas  is  set  free  from  the  anion 
Ag(CN)2',  and  the  anode  becomes  covered  with  silver  cyanide, 
which  soon  interrupts  the  current. 

All  the  silver  salts,  particularly  the  chloride,  bromide  and  iodide, 
are  sensitive  to  light,  i.e.  they  are  decomposed  by  light,  espe- 
cially by  the  violet  and  ultra-violet  rays  of  the  spectrum;  as  a  result, 
the  halogen  passes  off  and  the  color  of  the  salt  becomes  first  violet 
and  then  black.  A  blackened  preparation  of  this  sort  can  be  re- 
whitened  by  chlorine-or  bromine-water.  The  sensitiveness  to  light 
depends  in  large  measure  on  the  manner  in  which  the  silver  halide 
is  precipitated. 

247.  Photography.  The  property  of  silver  chloride  and  silver 
bromide  just  mentioned  forms  the  basis  of  photography.  The 
process  is  essentially  as  follows:  A  glass  plate  is  coated  with  a 
""sensitive  film,"  i.e.  a  thin  layer  of  silver  chloride  or  bromide  is 
spread  over  it.  Formerly  this  was  usually  prepared  by  the  pho- 
tographers themselves  from  collodion  (see  ORG.  CHEM.,  §  231) 
which  contained  a  halogen  salt,  e.g.  CdI2,  in  solution.  After 
the  evaporation  of  the  solvent  a  halide  coating  remained,  and  by 
•dipping  the  plates  so  prepared  into  a  solution  of  silver  nitrate, 
the  silver  halide  was  formed  on  them.  These  were  the  "wet 
plates";  now  they  are  almost  entirely  superseded  by  the  "dry 
plates." 

The  latter  are  prepared  commercially  on  a  large  scale.  They 
consist  of  a  film  of  silver  bromide  in  gelatine  (less  frequently  in 
collodion)  on  a  glass  plate. 

A  sensitive  plate  of  this  sort  is  placed  in  the  photographic 
apparatus,  which  is  essentially  a  camera  obscura,  and  the  plate  is 
there  "exposed"  to  a  light-image,  which  affects  the  silver  halide 
chemically.  It  is  very  probable  that  by  the  action  of  the  light  a 
subhalide  is  formed;  the  liberated  bromine  enters  into  combination 
with  the  gelatine  or  -the  collodion  and  is  therefore  unable  to  trans- 
form the  subhg^Jfc  into  halide.  As  yet  no  picture  can  be  seen  on 
the  plate;  it  must  first  be  "developed."  For  the  latter  purpose 
the  plate  is  immersed  in  a  liquid  containing  a  reducing  substance. 
A  typical  developer  is  a  solution  of  ferrous  oxalate  in  an  excess 
of  potassium  oxalate;  various  other  organic  compounds  (amido- 


247.]  PHOTOGRAPHY.  365 

phenols,  etc.)  are  at  present  frequently  used.  At  those  places  on 
the  plate  where  the  light  has  struck,  more  or  less  silver  (according 
to  the  intensity  of  the  action  of  the  light)  is  set  free  in  the  metallic 
form  as  a  very  thin  coating,  while  the  remaining  silver  halide  is  not 
affected  by  the  developer.  This  halide  must  next  be  removed,  else 
it  would  be  decomposed  by  the  light  and  more  silver  liberated; 
therefore  it  is  immersed  in  a  solution  of  sodium  thiosulphate 
("hypo")-  This  operation  is  called  "fixing"  the  image.  Up  to 
this  time  the  plate  must  be  kept  from  the  light. 

After  the  fixing  we  have  a  so-called  negative,  i.  e.  there 
remains  on  the  glass  plate  a  picture  which  is  black  in  those  places 
which  were  illuminated  in  the  object  and  clear  on  those  places  which 
were  dark.  From  this  a  positive  impression  is  prepared  by 
laying  the  negative  on  a  paper  coated  with  a  sensitive  film  and 
exposing  the  whole  to  direct  sunlight.  Those  places  on  the  nega- 
tive where  silver  was  deposited  let  no  light  or  very  little  through 
(according  to  their  thickness),  so  that  a  positive  image  is  now  pro- 
duced. Finally  the  positive  image  is  also  fixed,  for  which 
purpose  a  bath  containing  thiosulphate  and  a  little  gold  chloride 
is  used.  The  latter  improves  the  color- 1  o  n  e  of  the  photograph. 

The  photographic  process  in  its  various  stages  is  very  interesting  also 
from  a  theoretical  standpoint  and  deserves  a  little  more  detailed  study. 

1.  Preparation  of  the  Plates. — A  mixture  is  made  of  solutions  of  silver 
nitrate  and  ammonium  bromide  containing  enough  gelatine  to  make  them 
congeal  at  room  temperature.  No  separation  of  silver  bromide  is  observed 
immediately  on  mixing,  as  is  the  case  when  the  corresponding  aqueous 
solutions  are  mixed.  It  may  be  assumed  that  the  gelatine  acts  as  a  protective 
colloid  toward  the  silver  bromide,  which  of  itself  is  unable  to  form  a  hydrosol 
(cf.  §  196).  That  silver  bromide  is  really  formed  can  be  demonstrated 
by  measuring  the  electrical  conductivity.  If,  instead  of  the  silver  bromide, 
its  ions,  Ag'  and  Br',  were  present,  the  conductance  would  have  to  be  much 
greater  than  that  corresponding  to  the  ammonium  nitrate  which  results 
from  the  mixing  (AgNO3  +  NH4Br  =  AgBr  +  NH4NO3).  The  observed  con- 
ductance is,  however,  very  nearly  equal  to  that  of  a  gelatinous  solution 
of  ammonium  nitrate  having  the  same  concentration.  This  freshly  pre- 
pared cofloidal  silver  bromide  in  gelatine  is  relatively  not  very  sensitive  to 
light.  In  order  to  increase  its  sensitiveness  the  mixture  is  allowed  to  "ripen" 
by  standing  in  the  warm  for  a  considerable  length  of  time.  It  then  loses 
its  transparency  and  becomes  yellowish  white.  The  resulting  increase  in 
sensitiveness  must  be  accounted  for  by  supposing  that  the  light  is  not 
sufficiently  absorbed  by  the  transparent  colloidal  silver  bromide  to  exert 
its  full  action,  and  that  this  is  only  accomplished  when  in  the  process  of 
ripening  the  colloid  is  slowly  coagulated,  the  finer  particles  of  silver  bromide 


366  INORGANIC    CHEMISTRY.  [§§247- 

having  collected  to  form  larger  ones,  which  render  the  mass  opaque  and 
therefore  increase  its  absorptive  power.  The  ripened  silver  bromide  gela- 
tine is  then  spread  upon  the  plates. 

2.  The  Latent  Image. — When  the  plates  are  exposed  to  light  there  is 
formed  on  all  places  that  the  light  has  affected,  a  "  photohaloid,"  i.e.,    a 
mixture  of  silver  bromide  with  some  ultra-microscopic  particles  of  silver. 
When  AgBr  is  exposed  to  the  light  ordinarily  (apart  from  gelatine),  free 
bromine  is  formed  during  the  exposure ;  if  a  closed  apparatus  is  used  and 
it  is  afterward  placed  in  the  dark,  silver  bromide  is  formed  again.     More- 
over, not  all  the  silver  bromide  is  decomposed,  but  an  equilibrium  is 

established : 

AgBr?=±Ag+Br, 

which  is  displaced  farther  to  the  right,  the  stronger  the  illumination. 
Light  thus  plays  the  same  role  in  this  dissociation  as  heat  in  other  dis- 
sociations. If  a  gelatine  plate  is  used  the  latent  image  remains  for 
months  unaltered  because  the  free  bromine  is  taken  up  by  the  gelatine. 

3.  Developing. — This  process  is  explained  by  some  as  follows:    By 
the  reducing  action  of  the  developer  silver  is  immediately  set  free  from 
silver  subbromide  but  not  from  silver  bromide,  notwithstanding  that 
the  latter  is  capable  of  being  reduced.    The  system  silver  bromide  plus 
developer  can  be  compared  to  a  supersaturated  solution,  which  only 
deposits  solid  salt  when  it  comes  in  contact  with  a  crystalline  nucleus  of 
salt  (cf.  §  237).    The  nuclei  of  metallic  silver  are  furnished  by  the  silver 
separated  out  of  the  subbromide.    The  deposition  of  additional  silver 
molecules  takes  place  only  upon  those  molecules  already  there  and 
not  on  spots  where  there  was  no  subbromide  originally,  i.e.  silver  is 
deposited  only  where  the  light  acted  on  the  plate.     According  to  this 
nucleus  theory  the  developing  process  would  be  comparable  to  the 
following  experiment:  If  a  few  letters  are  written  on  a  glass  plate  with 
a  piece  of  alum  and  the  plate  is  laid  in  a  supersaturated  solution  of 
this  salt,  the  letters  become  visible,  because  alum  is  deposited  on  them. 

Silver  sulphate,  Ag2S04,  is  obtained  by  dissolving  silver  in 
hot  concentrated  sulphuric  acid.  It  is  scarcely  soluble  in  cold 
water. 

Silver  nitrate,  AgN03,  prepared  by  dissolving  silver  in  nitric 
acid,  crystaUizes  isomorphous  with  saltpetre  in  beautiful  rhombic 
crystals.  It  is  very  soluble  in  water  (1  part  in  0.5  part  at  room 
temperature)  and  melts  at  218°.  In  medicine  it  is  frequently 
employed,  especially  as  a  caustic;  it  goes  under  the  name  of  "lunar 
caustic."  Indelible  inks  are  also  prepared  from  it. 


248.]  GOLD.  367 

Silver  nitrite,  AgN(>2,  is  formed  as  a  yellowish  precipitate  on 
mixing  an  aqueous  alkali  nitrite  solution  and  silver  nitrite ;  it  dis- 
solves in  boiling  water  and  crystallizes  on  cooling  in  beautiful 
needles. 

GOLD. 

248.  This  metal  generally  occurs  native,  being  found  in  beds 
of  quartz  and  alluvial  deposits  resulting  from  the  decay  of  quartz 
rocks.  Traces  of  gold  have  been  detected  in  sea-water.  It  occurs 
in  Hungary,  Transylvania,  the  Ural  and  particularly  in  Australia, 
in  Transvaal  and  in  the  western  part  of  the  United  States  and 
Canada.  Recently  large  quantities  have  been  discovered  in  Alaska 
(Klondike  region) .  In  Colorado  considerable  gold  is  obtained  from 
tellurides  (sylvanite,  etc.). 

Inasmuch  as  the  amount  of  gold  contained  in  a  cubic  meter  of 
ore  or  rock  in  the  most  profitable  instances  is  only  very  small,  it 
becomes  the  task  of  metallurgy  to  extract  it  from  proportionately 
large  quantities  of  rock. 

In  the  Transvaal  this  is  accomplished  as  follows:  The  gold  occurs 
there  in  so-called  reefs,  which  are  vertical  veins  in  the  quartz.  These 
reefs  are  seldom  more  than  one  meter  thick,  but  extend  for  miles  east  and 
west;  their  depth  is  unknown.  They  are  mined  by  blasting  with  dyna- 
mite; the  large  pieces  are  reduced  to  about  the  size  of  an  egg  in  a  heavy 
iron  apparatus  and  then  sent  to  the  stamps,  that  move  in  a  large  trough 
through  which  plenty  of  water  is  kept  running.  The  water  carries  off 
the  fine  auriferous  slime,  which  is  made  to  flow  over  amalgamated  copper 
plates  that  are  somewhat  inclined.  The  gold  is  retained  by  the  mercury. 
After  some  time  the  plates  are  scraped  off  and  the  mercury  removed  by 
distillation,  leaving  the  gold. 

The  extracted  slime  ("tailings  ")  is  treated  again  for  gold,  for  which 
purpose  the  cyanide  process  of  SIEMENS  is  employed. 

By  this  process  the  tailings  are  allowed  to  stand  for  from  one  day  to 
three  weeks  in  contact  with  a  0.1  to  0.01%  potassium  cyanide  solution. 
Under  the  influence  of  the  oxygen  of  the  air  the  gold  dissolves  in  it, 
forming  a  double  cyanide,  KAu(CN)2: 

2Au +4KCN+ 2H2O+ O2=  2KAu(CN)2+  2KOH  +  H202. 

Hydrogen  peroxide  is  also  formed  and  serves  to  bring  further  amounts  of 
gold  into  solution : 

2Au+4KCN+HA  =  2KAu(CN)2+2KOH. 


368  INORGANIC  CHEMISTRY.  [§§  248- 

From  this  solution  the  gold  is  obtained  by  electrolysis  between  steel 
anodes  and  lead  cathodes.  At  the  anode  Prussian  blue  ( §  308)  is  formed, 
which  is  treated  for  potassium  cyanide;  the  gold  is  deposited  at  the 
cathode  (§  246).  This  gold  is  separated  from  the  lead  it  contains  by 
cupellation. 

[Particularly  in  the  United  States  two  processes  (chlorination  and 
cyanide)  are  in  general  use  for  extracting  gold  from  its  ores  without 
amalgamation.  Both  processes  are  especially  applicable  to  low-grade 
and  sulphurous  ores,  e.g.  the  tellurides  of  Colorado.  In  the  chlorination 
process  the  ore  is  crushed  and  roasted  and  then  treated  in  revolving 
barrels  with  chlorine,  prepared  either  chemically  or  electrolytically, 
after  which  the  gold  is  precipitated  with  hydrogen  sulphide  and  roasted. 
The  cyanide  process  is  much  similar  to  that  described  above  for  treating 
the  tailings,  but  zinc  generally  serves  as  the  precipitant  instead  of  elec- 
trolysis. 

Placer  and  hydraulic  mining  find  application  in  newly  discovered 
deposits  but  are  much  less  common  than  vein  mining.  For  the  present 
status  of  the  metallurgy  of  gold  as  well  as  other  metals  the  student 
should  consult  a  mining  annual. — TR.] 

249.  Physical  Properties. — When  pure,  gold  is  reddish  yellow, 
very  soft,  (much  like  lead)  and  extremely  malleable  and  ductile. 
The  thinnest  gold-leaf  appears  green  in  transmitted  light.  Sp. 
g.  =  19.265  at  13°.  It  is  a  very  good  conductor  of  heat  and  elec- 
tricity. At  1063°  it  melts  to  a  greenish  liquid. 

Chemical  Properties. — Gold  is  the  typical  representative  of  the 
precious  metals;  it  is  not  attacked  by  acids  and  is  dissolved  only 
by  chlorine-water,  aqua  regia  and  potassium  cyanide  solution  (see 
above).  Its  compounds  are  all  very  unstable;  on  warming  they 
decompose,  leaving  the  metal. 

Uses. — About  one-half  the  world's  production  of  gold  is  used 
for  industrial  purposes.  For  these  purposes  the  pure  metal  is  too 
soft,  however,  and  must  be  alloyed  with  copper  or  silver.  The 
proportion  of  gold  in  the  alloy  is  ordinarily  expressed  in  carats; 
the  pure  metal  is  24  carats;  gold  jewelry,  etc.,  usually  14-18  carats 
i.e.,  24  parts  of  the  alloy  contain  14-18  parts  of  gold.  v  The  gold 
coins  of  the  United  States  contain  1  part  copper  to  9  parts  gold, 
those  of  England  1  part  copper  to  11  parts  gold. 

For  purposes  of  gold-plating  the  same  electrolytic  processes  are 
employed  as  for  silver-plating. 


249.1  TESTING    OF    GOLD    AND    SILVER.  369 


Testing  of  Gold  and  Silver. 

The  oldest  method  of  testing  is  by  means  of  the  touchstone,  or  '  '  Lydian 
stone/'  a  black  basalt.  This  stone  must  be  dull  black,  unaffected  by 
aqua  regia  and  somewhat  rough.  The  sample  is  rubbed  on  the  surface 
of  the  stone  so  as  to  leave  a  bright  streak  of  particles  of  the  metal.  This 
streak  is  then  compared  with  that  of  a  series  of  touchneedles  of  known 
composition. 

Silver  streaks  are  compared  merely  as  to  color.  A  skilled  observer 
can  usually  estimate  the  proportion  of  silver  to  within  2.0-1.5%. 

In  the  case  of  gold  objects  it  must  be  known  whether  the  metal  con- 
tains copper,  silver,  or  both.  Therefore  the  color  of  the  streak  is  com- 
pared with  that  of  touchneedles  of  the  presumably  corresponding  alloy. 
The  streaks  are  then  moistened  with  a  little  acid  consisting  of  1  part  HC1, 
SO  HX03  and  100  H20.  Alloys  with  75%  or  more  gold  are  not  attacked 
by  this  mixture  at  ordinary  temperatures.  If  the  percentage  is  less,  it  is 
possible  to  detect  differences  of  1%.  This  method  is  decidedly  crude 
and  is  usually  employed  only  in  confirming  a  supposed  percentage. 
Where  the  metal  is  rich  it  is  very  deceptive;  but  gold  of  a  quality  such 
as  is  generally  used  for  ornaments,  etc.  (ca,  I5o8o3<r  fine)  can  be  safely  tested 
in  this  way. 

A  far  more  reliable  test  of  the  quality  of  gold  is  by  cupellation.  Part 
of  the  sample  is  fused  with  lead  in  a  small  muffle  furnace  in  a  small  thick- 
walled,  porous  crucible  (cupel)  consisting  of  bone  ash.  At  the  high  tem- 
perature of  the  furnace  the  lead  and  any  copper  present  are  oxidized  and 
their  oxides  melt  and  are  absorbed  by  the  bone-ash.  As  soon  as  they  are 
completely  taken  up  the  brilliant  surface  of  the  metal  suddenly  appears 
with  splendid  effect  (brightening  of  gold).  The  residual  drop  of  metal  is 
an  alloy  of  only  gold  and  some  silver.  The  latter  is  got  rid  of  with 
boiling  nitric  acid,  after  the  alloy  left  on  the  cupel  has  solidified  and  been 
hammered  flat. 

The  proportion  of  silver  in  silverware  is  now  determined  exclusively  in 
the  wet  way,  by  titration. 

250.  Gold  forms  two  series  of  compounds  analogous  to  the 
oxides  Aii2O,  aurous  oxide,  and  Ai^Oa,  auric  oxide. 


Aurous  Compounds. 

Aurous  oxide,  Au20,  is  obtained  by  treating  the  aurous  chloride 
with  dilute  potassium  hydroxide.  It  is  a  dark  violet  powder, 
which  breaks  up  into  its  constituents  at  250°. 

Aurous  chloride,  AuCl,  is  produced  by  heating  auric  chloride 
to  185°.  It  is  white  and  insoluble  in  water.  When  heated  it 


370  INORGANIC  CHEMISTRY.  [§§250- 

splits  up  into  its  elements.  On  being  warmed  with  water  it  yields 
2Au  and  AuCl3. 

Aurous  iodide  is  formed  (like  cuprous  iodide)  on  treating  a 
solution  of  the  chloride  with  potassium  iodide. 

The  gold  double  cyanide,  KCN-AuCN,  is  prepared  by  dissolving 
auric  oxide  in  potassium  cyanide;  it  is  used  in  gold-plating. 

Of  the  oxy-salts  of  aurous  oxide  only  a  few  double  salts  are 
known. 


Auric  Compounds. 

Auric  chloride,  AuCl3,  can  be  obtained  by  dissolving  gold  in 
aqua  regia  or  by  the  action  of  chlorine  on  the  metal.  It  forms  a 
dark  red  crystalline  mass,  which  deliquesces  rapidly.  On  the 
evaporation  of  its  solution  it  partially  decomposes  into  chlorine 
and  aurous  chloride.  By  evaporating  with  hydrochloric  acid  long 
yellow  needles  are  obtained,  consisting  of  a  compound  AuCl3»HCl, 
which  can  be  regarded  as  chlor-auric  acid.  Many  salts  of  this  acid 
are  known  to  exist,  e.g.  KC1  - AuCl3  +  2 JH20  and  NH4Cl.AuCl3  + 
H20,  as  well  as  many  chlor-aurates  of  organic  bases.  These  double 
salts  give  the  ordinary  tests  for  gold,  hence  this  acid  either  forms 
no  complex  ion  AuCU'  or  is  very  unstable.  Auric  chloride  is  also 
soluble  in  alcohol  and  in  ether. 

Auric  oxide,  Au203,  can  be  prepared  by  precipitating  auric 
chloride  with  magnesia.  The  latter  can  be  removed  from  the  pre- 
cipitate with  concentrated  nitric  acid,  the  auric  oxide  remaining 
as  a  brown  powder,  which  breaks  up  at  250°  into  its  elements. 

If  the  precipitate  produced  by  magnesia  is  treated  with  dilute 
nitric  acid,  a  reddish-yellow  powder  of  the  formula  Au03H3  is 
obtained,  which  displays  acid,  instead  of  basic,  properties.  Salts 
of  this  auric  acid  are  known,  which  are  derived  from  the  com- 
pound Au(OH)3  — H20=AuO-OH.  Potassium  aurate,  for  exam- 
ple, has  the  formula  KAu02+3H20  and  crystallizes  in  yellow 
needles.  Many  other  salts  are  also  known;  the  above-mentioned 
precipitate  with  magnesia,  for  example,  can  be  looked  upon  as  the 
magnesium  salt  of  auric  acid,  Mg(Au02)2. 

Auric  sulphide,  Au2S3,  is  precipitated  from  gold  solutions  by 
hydrogen  sulphide.  It  is  very  dark  brown  and  soluble  in  am- 
monium sulphide. 

Gold  is  precipitated  from  its  solutions  in  the  metallic  form  by 


251.]  AURIC  COMPOUNDS.  371 

various  reducing  agents.  Ferrous  sulphate  (§  248),  oxalic  acid 
and  acetylene  water  serve  very  well  for  this  purpose.  Hydrogen 
peroxide  precipitates  gold  quickly  in  alkaline  solution. 

251.  For  many  centuries  the  alchemists  endeavored  to  produce  gold 
from  the  baser  metals.  It  is  needless  "to  say  that  their  efforts  were  never 
rewarded.  The  chances  of  this  hope  being  realized  must  at  present  be 
regarded  as  very  slight,  since  gold  is  an  element.  Inasmuch,  however,  as 
our  conception  of  an  element  is  relative  (§8),  i.e.  it  depends  on  the 
extent  of  our  mastery  over  natural  forces,  the  impossibility  of  decom- 
posing gold  or  synthesizing  it  from  other  elements  >s  by  no  means  abso- 
lutely established.  (See  §  §  266-7) . 

Although  we  now  ascribe  to  every  metal  fixed,  unalterable  properties, 
it  might  well  have  seemed  possible  to  the  alchemists,  with  their  more 
limited  knowledge,  that  the  properties  of  the  metal?  could  vary.  None 
of  the  metals  except  gold  occur  pure  in  nature;  they  have  to  be  extracted 
from  oxides  or  sulphides,  which  frequently  contain  various  impurities. 
The  metals  thus  obtained  had  no  definite  properties;  distinction  was 
made  between  various  sorts  of  lead,  copper,  etc.  The  mutability  of 
the  metals  may  be  said  to  have  been  the  first  principle  which  observation 
taught;  indeed,  when  a  piece  of  metal  is  fused  with  small  amounts  of 
various  other  substances,  its  properties  (color,  etc.)  really  do  change. 
Moreover,  at  the  time  of  the  alchemists  the  present  concept  "element  " 
was  not  yet  established;  this  was  first  introduced  by  BOYJ-E  (1627-1691). 
Before  then,  the  doctrine  of  ARISTOTLE  was  very  generally  accepted, 
according  to  which  all  substances  are  made  up  of  air,  fire,  earth  and 
water.  In  order  to  produce  gold  it  therefore  seemed  only  necessary  to 
deprive  the  baser  metals  of  certain  properties  and  substitute  others.  As 
to  the  metals  themselves  the  idea  was  prevalent  in  alchemistic  circles 
that  mercury  was  the  primordial  substance  and  that  it  had  undergone 
various  changes.  Before  gold  could  be  made  from  it  it  must  be  made 
refractory  and  of  a  yellow  color.  Not  a  few  alchemists  were  convinced, 
moreover,  that  the  success  of  the  "great  work  "  depended  on  the  ?ooper- 
ation  of  a  higher  power. 

• 

SUMMARY  OF  THE  GROUP. 

252.  The*  metals  copper,  silver  and  gold  form  a  bridge  from 
the  difficultly  fusible  metals,  Ni,  Pd,  Pt  (Group  VIII),  tc  the 
easily  fusible,  Zn,  Cd,  Hg  (Group  II);  their  melting-points  are 
between  those  of  the  two  groups.  The  following  brief  table  sum- 
marizes the  physical  constants  of  these  metals  as  well  as  those*  of 
the  related  elements,  lithium  and  sodium: 


372 


INORGANIC  CHEMISTRY, 


[§§  252- 


Li 

Na 

Cu 

Ag 

Au 

Atomic  weight  

7.00 

23.00 

63.57 

107.88 

197.2 

Specific  gravity          •  • 

0  59 

0.97 

8.94 

10  5 

19.33 

Melting-point      

180 

97.6 

1083 

961 

1063 

Color            

white 

white 

red 

white 

red 

The  analogy  in  the  chemical  properties  is  chiefly  apparent  in 
the  -ous  compounds.  These  have  the  type  R20  for  the  oxygen 
compounds  and  RX  for  the  halides.  The  -ous  halides  of  Cu,  Ag 
and  Au  are  all  white  and  insoluble  in  water;  they  are  isomorphous 
with  sodium  chloride. 

Moreover,  there  are  certain  analogies  in  solubility.  Lithium 
carbonate  and  hydroxide  are  less  soluble  in  water  than  the  corre- 
sponding sodium  compounds;  copper  carbonate  and  hydroxide 
are  insoluble,  while  the  corresponding  silver  compounds  dissolve 
to  some  extent.  The  sulphate  of  sodium  (third  horizontal  series) 
crystallizes  preferably  with  10H20,  that  of  copper  (fifth  series) 
with  5H2O,  while  silver  sulphate  (seventh  series)  is  anhydrous. 

The  oxygen  compounds  exhibit  a  gradual  decrease  in  stability. 
Li2O  and  Na2O  are  unaffected  by  high  temperatures,  but  CuO  is 
transformed  into  Cu20,  and  the  oxides  of  silver  and  gold  break  up 
even  at  comparatively  low  temperatures  into  their  elements. 

However,  it  must  be  admitted  that  the  analogy  between  these 
elements  is  not  so  great  as  in  other  groups.  Their  difference  in 
valence  is  especially  striking  and,  moreover,  there  is  little  simi- 
larity in  the  properties  of  the  higher  stages  of  oxidation.  This  is 
one  of  the  weak  parts  of  the  periodic  system. 


BERYLLIUM     AND    MAGNESIUM. 
I.  Beryllium  (Glucinum). 

253.  This  is  one  of  the  rarer  elements.  It  occurs  in  the  mineral  beryl, 
Al2O3-3Si02  +  3(BeO-SiO2);  that  variety  of  beryl  which  is  colored  green 
by  traces  of  a  chromium  compound  is  the  gem  called  emerald,  or  smaragd. 
Chrysoberyl  has  the  composition  BeO-Al203. 

Almost  all  the  beryllium  compounds  are  made  from  beryl.  This  is 
disintegrated  by  fusing  with  potassium  carbonate.  The  fused  mass,  after 
cooling,  is  treated  with  sulphuric  acid  to  precipitate  the  silica.  Most  of 
the  aluminium  is  then  removed  by  crystallization  in  the  form  of  alum, 


253.]  BERYLLIUM  AND  MAGNESIUM.  373 

as  this  is  sparingly  soluble  in  cold  water,  while  beryllium  sulphate  remains 
in  the  mother  liquor.  The  latter  is  then  mixed  with  a  hot  solution  of 
ammonium  carbonate  to  precipitate  aluminium  and  iron,  beryllium  still 
remaining  in  solution.  After  acidifying  with  hydrochloric  acid,  the 
beryllium  is  precipitated  as  the  hydroxide  by  ammonia 

The  metal  was  obtained  by  heating  the  double  fluoride  BeF2-2KF 
with  sodium.  It  is  a  malleable  solid  with  the  specific  gravity  1.64.  It 
does  not  decompose  water,  even  at  100°.  At  ordinary  temperatures 
it  is  permanent  in  the  air.  Hydrochloric  and  sulphuric  acids  dissolve 
it  readily  with  the  evolution  of  hydrogen;  dilute  nitric  acid  does  not 
attack  it  so  readily.  Beryllium  is  also  dissolved  easily  by  caustic  potash 
and  soda  with  the  evolution  of  hydrogen  and  the  formation  of  salts 
having  the  formula  Be(OR)2.  The  hydroxide  thus  behaves  as  a  weak 
acid  towards  strong  bases.  These  properties  correspond  to  those  of 
aluminium;  in  §  215  attention  was  already  called  to  the  analogy  between 
these  two  elements.  This  analogy  also  characterizes  their  compounds, 
e.g.,  beryllium  carbide  yields  pure  methane  with  water,  just  like  alu- 
minium carbide  (§178). 

Only  one  oxide  of  beryllium  is  known,  BeO  (§  215).  It  is  a  white 
powder,  which  after  ignition  is  difficultly  soluble  in  acids  (like  A1203). 
It  is  obtained  by  heating  the  hydroxide,  Be(OH)2,  which  is  precipi- 
tated from  solutions  of  the  salts  as  a  white  gelatinous  mass.  When 
freshly  precipitated,  it  is  easily  soluble  in  alkalies,  ammonium  carbonate 
and  dilute  acids.  On  being  heated  with  water,  dilute  ammonia  solution 
or  dilute  alkali  solution,  or  on  being  ignited,  or  even  on  standing  for  some 
time,  it  "grows  old"  and  loses  these  properties.  Heating  with  ten-fold 
normal  solution  of  an  alkali  hydroxide  " rejuvenates"  even  the  "  oldest" 
beryllium  hydroxides,  which  are  dissolved  only  slowly  by  warm  concen- 
trated hydrochloric  acid.  Beryllium  hydroxide  is  distinguished  from 
aluminium  hydroxide  in  two  respects:  it  dissolves  in  ammonium  carbonate 
(see  above)  and  is  precipitated  from  the  solution  in  caustic  soda  or 
caustic  potash  by  prolonged  boiling. 

Beryllium  sulphate,  BeSO4  crystallizes  with  four  or  seven  molecules  of  water, 
in  the  latter  case  being  isomorphous  with  MgSO4-7H2O.  The  double  salt 
BeSO4-K2SO4-3H2O  is  (like  alum)  sparingly  soluble  in  cold  water.  Beryllium 
chloride,  BeCl2,  must  be  prepared  from  the  oxide  by  heating  with  charcoal  in 
a  current  of  chlorine.  Its  vapor  density  corresponds  to  the  formula  BeCl2.  It 
crystallizes  with  4H2O.  Beryllium  carbonate  is  soluble  in  water.  It  loses 
carbon  dioxide  very  easily. 

The  beryllium  salts  taste  sweet,  hence  the  name  g  1  u  c  i  n  u  m  (or  glycin- 
ium),  which  is  common  in  France  and  America. 


374  INORGANIC  CHEMISTRY.  [§§  254- 

II.  Magnesium. 

254.  This  element  occurs  as  carbonate,  silicate,  and  chlo- 
ride in  considerable  quantities.  Magnesite  is  MgC03,  dolomite 
MgCa(C03)2.  Among  the  silicates  containing  magnesium  we  have 
talc  and  soapstone,  H2Mg3Si4Oi2;  serpentine  (asbestos),  H4Mg3Si209; 
meerscliaum,  H4Mg2Si30i0.  It  is  found  in  smaller  amounts  in 
many  other  silicates,  e.g.  hornblende  (asbestos) ,  augite,  tourmaline. 
Other  salts  found  in  nature  are  carnallite,  MgCl2  •  KC1  •  6H20,  kie- 
serite,  MgS04-H2O,  and  kainite,  MgS04-KCl-3H2O  (Stassfurt 
Abraum  salts).  Upon  the  weathering  of  the  silicates  the  mag- 
nesium goes  into  the  soil,  whence  it  is  absorbed  by  the  plants  (to 
which  this  element  is  invaluable)  and  finally  taken  into  the  animal 
body. 

The  metal  is  manufactured  on  a  large  scale,  since  it  is 
employed  for  illumination  in  photography,  pyrotechnics,  etc.,  on 
account  of  the  intense  light  (flash-light)  produced  by  its  combustion. 
At  present  it  is  prepared  mainly  by  the  electrolysis  of  fused  mag- 
nesium chloride  or  carnallite  in  a  cast-steel  crucible,  which  serves 
as  cathode;  gas  carbon  is  used  for  the  anode.  It  is  also  obtained 
by  heating  the  double  chloride  MgCl2-NaCl  with  sodium.  It  is 
silvery-white  and  has  a  high  lustre.  Sp.  g.  =1.75.  It  is  malleable 
and  ductile  and  comes  on  the  market  in  the  form  of  wire  or  ribbon 
as  well  as  powder,  but  the  ribbon  frequently  contains  zinc.  It 
melts  at  651°  and  boils  at  1120°.  It  is  quite  permanent  in  the 
air,  since  it  soons  becomes  coated  with  a  thin  cohesive  film  of 
the  oxide;  at  an  elevated  temperature  it  burns  to  magnesia,  MgO. 
When  it  is  heated  red-hot  in  a  limited  supply  of  air,  a  large  part 
is  converted  into  the  nitrite,  Mg3N2,  a  yellowish-green  substance. 
Boiling  water  decomposes  it  slowly  with  the  evolution  of  hydrogen. 
It  dissolves  readily  in  acids  but  is  unaffected  by  alkalies.  It  is  a 
powerful  reducing-agent,  reducing  silica  (§  190),  for  example; 
moreover,  when  ignited,  it  burns  in  water  vapor. 

Magnesium  oxide,  MgO,  magnesia,  is  the  only  oxide  of  mag- 
nesium known.  It  results  from  the  combustion  of  the  metal  or 
from  heating  the  hydroxide  or  carbonate.  It  is  a  white,  very 
light  powder,  which  is  employed  in  medicine  under  the  name  mag- 
nesia usta.  With  water  it  forms  the  hydroxide  Mg(OH)2. 

Magnesium  hydroxide,  Mg(OH)2,  is  precipitated  from  solutions 
of  magnesium  salts  by  alkalies.  It  is  slightly  soluble  in  water  and 


255.]  MAGNESIUM  SALTS.  375 

turns  red  litmus  blue;  however  in  an  excess  of  alkali  its  ionization 
is  so  diminished  that  it  becomes  practically  insoluble.  It  is  a 
weak  base,  but  is  strong  enough  to  absorb  carbon  dioxide  from 
the  air.  It  dissolves  readily  in  an  aqueous  solution  containing 
ammonium  salts.  According  to  OSTWALD,  this  is  to  be  explained 
as  follows:  The  solution  of  an  ammonium  salt  contains  a  large 
quantity  of  NH4-ions.  When  a  substance  is  introduced  into  the 
solution,  which  gives  off  OH-ions;  as  does  magnesium  hydroxide, 
these  NH4-ions  unite  with  OH-ions  to  form  NH4OH,  or  rather 
NH3+H20  (cf.  §  234).  As  a  result  of  this  reaction  OH-ions  dis- 
appear. In  order  to  restore  the  equilibrium  between  the  undis- 
solvecl  magnesium  hydroxide  and  the  solution,  more  of  this  hy- 
droxide must  go  in  solution,  but  again  the  freshly  formed  OH-ions 
are  taken  up  by  the  NH4-ions.  If  sufficient  of  the  latter  are  present, 
this  process  will  go  on  till  all  the  magnesium  hydroxide  has  entered 
into  solution.  It  now  becomes  clear  why,  on  the  other  hand,  the 
solution  of  a  magnesium  salt  is  not  precipitated  by  ammonia  in 
the  presence  of  a  sufficient  quantity  of  ammonium  salt. 

MAGNESIUM    SALTS. 

255.  Magnesium  chloride,  MgCl2,  crystallizes  with  six  mole- 
cules of  water  and  is  very  hygroscopic.  The  deliquescence  of 
common  salt  is  due  to  the  magnesium  salt  it  usually  contains. 

OTT 

On  evaporating  the  aqueous  solution  the  basic  chloride,  Mgnl    , 

01 

and  hydrochloric  acid  are  formed;  sea-water  cannot  be  used  in 
boilers  because  of  the  magnesium  salt  it  contains,  for  the  hydro- 
chloric acid  set  free  attacks  the  iron.  Many  double  salts  of  mag- 
nesium chloride  are  known. 

It  can  be  obtained  anhydrous  by  heating  the  double  chloride 
MgCl2  •  NH4C1  •  6H2O,  when  it  forms  a  laminar-crystalline  mass, 
which  melts  at  708°  and  distils  without  decomposition  at  bright 
red  heat. 

Careful  study  of  the  decomposition  of  magnesium  chloride  by  oxygen 
and  by  steam  has  shown  that  a  reversible  reaction  is  involved  in  each 


case: 


2MgCl2  +  02z=»2MgO  +2C12;      MgCl2  +  H2O<=±MgO  +2HC1. 


In  the  former  reaction  a  rise  of  temperature  displaces  the  equilibrium 
toward  the  right,  although  below   500°  the   velocity  is  still  very  small. 


376  INORGANIC  CHEMISTRY.  [§§  255- 

In  the  second  process  the  composition  of  the  gaseous  equilibrium  mix- 
ture at  700°  has  been  found  to  be  90%  HC1  +  10%  H20. 

Magnesium  sulphate,  MgS04  •  7H20,  Epsom  salt,  finds  use  in 
medicine.  It  is  very  soluble  in  water.  It  loses  6  mols.  H2O  at 
150°,  and  the  seventh  above  200°.  In  this  respect  it  behaves  like 
other  sulphates,  e.g.  ZnS04-7H20,  FeS04-7H20,  and  those  of 
nickel  and  cobalt,  which  are,  moreover,  isomorphous  with  it.  A 
further  analogy  between  these  sulphates  appears  in  the  fact  that 
with  sulphate  of  potassium  or  ammonium  they  form  double  salts  of 
the  same  type,  K2S04-MgSO4-6H20,  which  are  also  isomorphous. 

Magnesium  ammonium  phosphates,  MgNH4P04  •  6H20,  serves 
for  the  precipitation  of  magnesium  as  well  as  of  phosphoric  acid. 
It  is  not  wholly  insoluble  in  water,  but  does  not  dissolve  in 
ammonia,  the  reason  for  which  is  again  to  be  found  in  the  reduc- 
tion of  the  ionization.  Completely  analogous  to  this  compound 
is  the  corresponding  arsenate,  MgNH4As04  •  6H2O. 

Magnesium  carbonate. — From  solutions  of  magnesium  salts 
soda  precipitates  a  basic  carbonate,  Mg(OH)2-4MgC03-4H20. 
The  carbon  dioxide  liberated  holds  part  of  the  magnesium  in  solu- 
tion as  acid  carbonate.  This  precipitate  is  known  as  magnesia 
alba.  The  neutral  carbonate  can  be  prepared  from  it  by  sus- 
pending magnesia  alba  in  water,  passing  in  carbon  dioxide  and 
allowing  to  stand;  in  time  the  salt  MgCO3-3H2O  crystallizes  out, 
which  is,  however,  readily  split  up  hydrolytically  by  water,  form- 
ing basic  carbonate  again. 

CALCIUM,  STRONTIUM  AND  BARIUM. 
I.  Calcium. 

256.  This  element  is  one  of  the  ten  principal  constituents  of 
the  earth's  crust  (§8).  Particularly  the  carbonate  is  found 
in  large  quantities  in  nature,  limestone,  calcite,  aragonite,  marble 
and  chalk,  all  being  forms  of  it.  An  earthy  deposit  containing  a 
certain  amount  of  calcium  carbonate  is  termed  marl.  Calcium 
silicates  and  especially  calcium  double  salts  constitute  the  major 
portion  of  the  siliceous  rocks.  There  are  also  extensive  beds  of 
calcium  phosphate,  phosphorite,  apatite,  etc.,  particularly  in  Spain 
and  Florida.  Calcium  occurs  as  sulphate  in  the  form 'of  gypsum 
and  alabaster.  Moreover,  in  the  animal  kingdom  large  quantities 


257.]  OXIDES  AND  HYDROXIDES  OF  CALCIUM.  377 

of  this  element  are  found.  The  skeletons  of  vertebrates  are  chiefly 
phosphate  and  carbonate  of  calcium;  the  shells  of  mollusks  con- 
sist of  calcium  carbonate,  as  do  also  eggshells.  As  for  the  plants, 
lime  is  one  of  their  indispensable  inorganic  constituents. 

Metallic  calcium  can  be  obtained  by  electrolysis  of  a 
fused  mixture  of  calcium  chloride  and  calcium  fluoride.  Such 
a  mixture  melts  much  lower  than  the  single  salts  (§  237).  The 
lower  temperature  makes  the  separation  of  the  metal  easier  and 
prevents  its  combustion.  Calcium  is  a  silvery-white  metal, 
which  melts  at  800°;  it  is  soft  enough  to  cut  and  is  malleable, 
but  less  so  than  potassium  and  sodium;  it  has  a  crystalline  frac- 
ture. Sp.  g.=  1.52.  It  is  relatively  little  affected  by  oxy- 
gen, chlorine,  bromine,  and  iodine,  all  of  which  react  with  the 
metal  only  at  a  higher  temperature  than  the  ordinary  one.  In 
a  current  of  air  calcium  unites  with  both  oxygen  and  nitrogen 
(§  110).  With  hydrogen  it  forms  a  compound  CaH2,  which  is 
also  prepared  commercially  by  passing  hydrogen  into  molten 
calcium.  The  calcium  hydride  reacts  with  water  most  vigorously 


Since  1  kilo  of  the  hydride  evolves  about  1  cubic  meter  of 
hydrogen,  it  constitutes  a  very  suitable  material  for  generating 
hydrogen  for  aeronautic  purposes,  especially  in  out-of-the-way 
places. 

OXIDES    AND    HYDROXIDES  OF   CALCIUM. 

257.  Calcium  oxide,  CaO,  (quick-lime,  unslaked  lime)  is  pre- 
pared commercially  by  "  burning  "  limestone  or  mollusk  shells. 
The  limestone  is  mixed  with  coal  and  the  latter  is  set  on  fire;  the 
heat  of  the  burning  coal  decomposes  the  carbonate  of  lime  into 
calcium  oxide  and  carbon  dioxide.  The  kilns  are  usually  con- 
structed in  such  a  way  that  the  burned  lime  can  be  drawn  out  at 
the  bottom  while  the  mixture  of  fuel  and  limestone  is  fed  in  at  the 
top,  so  that  the  process  is  continuous.  In  the  United  States 
"  long-flame  "  periodic  kilns  are  generally  used  because  they 
are  simpler  and  fuel  is  inexpensive. 

Calcium  oxide  is  a  white  amorphous  powder,  which  requires 
the  temperature  of  t^e  electric  r  r°  fli  mace  for  fusion  (§  176).  On 


373  INORGANIC  CHEMISTRY.  [§§257- 

being  heated  strongly  with  an  oxy-hydrogen  flame  it  emits  an 
intense  white  light  (§  13).  It  absorbs  water  and  carbon  dioxide 
from  the  air;  as  a  result  the  chunks  of  lime,  which  are  hard  and 
solid  when  they  come  from  the  kiln,  gradually  crumble  to  fine 
powder. 

Calcium  hydroxide,  Ca(OH)2,  (slaked  lime)  is  obtained  by 
"  slaking  "  quick-lime  with  water.  Its  formation  is  attended  by 
the  evolution  of  much  heat.  It  is  only  sparingly  soluble  in  water 
(forming  lime-water),  but  more  soluble  in  cold  water  than  in  warm. 
The  solubility,  is  however,  sufficient  to  make  the  precipitation  of 
this  hydroxide  by  ammonium  hydroxide  impossible,  for  the  con- 
centration of  the  hydroxyl  ions  of  the  latter  is  too  small  together 
with  that  of  the  calcium  ions  present  to  reach  the  value  of  the 
solubility  product  of  calcium  hydroxide.  At  red-heat  it  is  recon- 
verted into  the  oxide. 

Mortar. — Calcium  hydroxide  is  used  in  masonry.  For  this 
purpose  quicklime  is  mixed  with  water  and  sand  so  as  to  form  a 
thick  paste,  called  mortar,  which  is  thrown  in  between  the 
stones.  After  some  time  the  mass  becomes  as  hard  as  stone;  this 
is  due  to  the  conversion  of  the  hydroxide  into  the  carbonate  by 
the  action  of  the  carbon  dioxide  of  the  air.  The  sand  makes  the 
mass  porous,  so  that  the  process  of  hardening  extends  inward; 
the  older  the  wall  the  harder  the  mortar.  The  formation  of  cal- 
cium silicate  appears  to  play  only  a  minor  role  in  this  process. 

If  the  lime  contains  more  or  less  magnesia  it  is  difficult  to 
slake;  it  is  therefore  less  adapted  to  masonry  purposes  and  is 
called  "  poor,"  or  "  lean,"  in  contrast  with  tr^e  pure,  easily  slaked 
"  fat  "  lime. 

Cement  contains,  besides  lime  (50-60%),  principally  silica 
(ca.  24%)  and  alumina  (ca.  8%).  It  is  made  by  burning  a  mixture 
of  limestone,  clay  and  sand.  In  some  places,  e.g.  Brohlthal  in 
the  Rhine  region,  such  a  mixture  occurs  as  "  tuffstone,"  which 
yields  cement  directly  on  burning.  Cement  after  being  mixed 
with  water  sets  very  firmly  in  a  short  time;  this  is  due,  in  all 
probability,  to  the  fact  that  on  treating  it  with  water  calcium 
aluminate  is  dissolved  and  the  solution  slowly  deposits  a  hydrous 
aluminate,  which  is  much  less  soluble  and  causes  the  setting  of 
the  cement.  At  the  same  time  insoluble  calcium  aluminium 
silicates  are  formed. 


258.]  SALTS  OF   CALCIUM.  379 

Calcium  peroxide,  CaO2-SH2O,  is  deposited  when  lime-water 
is  treated  with  hydrogen  peroxide  solution.  It  gives  up  oxygen 
on  heating. 

SALTS    OF  CALCIUM. 

258.  Calcium  chloride,  CaCl2,  is  obtained  by  dissolving  the 
hydroxide  or  carbonate  in  hydrochloric  acid.  It  can  crystallize 
with  various  amounts  of  water.  The  hydrate  CaCl2-6H20  forms 
large  crystals.  Calcium  chloride  is  very  hygroscopic  and  is  there- 
fore frequently  used  to  dry  gases  or  to  absorb  water  dissolved  in 
organic  liquids  (ether,  carbon  disulphide,  etc.).  It  melts  at  719°. 
It  unites  with  ammonia  to  form  CaCl2-8NH3;  hence  it  cannot  be 
used 'to  dry  this  gas. 

When  crystallized  calcium  chloride  is  mixed  with  ice  the  temperature 
falls  considerably,  even  reaching  —48.5°.  Such  a  mixture  is  called  a 
cooling-  or  freezing-mixture  and  is  often  employed  for  producing  low 
temperatures.  Besides  calcium  chloride  and  ice,  many  other  such 
mixtures  are  known;  the  one  most  frequently  used  is  that  of  common 
salt  and  ice,  with  which  a  temperature  of  —21°  can  be  obtained.  Ice 
is  not  absolutely  necessary;  for  instance,  if  solid  ammonium  nitrate 
is  added  to  its  own  weight  of  water,  a  temperature  of  —15.5°  can  be 
produced. 

In  order  to  understand  why  such  mixtures  become  so  cold  we  must 
recall  §  237.  Suppose  that  ice  is  introduced  into  a  saturated  salt  solu- 
tion of  0°,  solid  s$lt  being  present  at  the  bottom  so  that  the  liquid 
remains  saturated.  The  system  solution  +  ice  is  not  in.  a  state  of  equi- 
librium at  0°,  for  the  salt  solution  has  a  freezing-point  much  lower  than  0°. 
It  cannot  therefore  continue  in  this  state,  but,  if  it  is  to  be  in  equilibrium 
with  ice  as  solid  phase,  the  temperature  must  sink,  and  this  is  only 
possible  as  the  ice  melts,  by  which  process  heat  is  changed  into  the 
latent  condition.  If  enough  ice  is  present,  it  can,  by  melting,  continue 
to  withdraw  free  heat  from  the  system  till  the  cryohydric  point  is 
reached;  for  only  at  or  below  that  point  can  ice  and  salt  exist  perma- 
nently side  by  side.  It  follows,  therefore,  that  the  cryohydric  tempera- 
ture is  the  lowest  that  can  be  reached  by  the  mixture.  In  §  237  it 
was  shown,  further,  that  there  is  no  essential  difference  between  the 
two  components  of  a  solution;  this  is  also  seen  on  considering  cooling- 
mixtures  containing  no  ice.  For  instance,  when  ammonium  nitrate  is 
added  to  water,  the  solution  has  a  freezing-point  much  lower  than  0°» 
Here  it  is  the  great  absorption  of  heat  in  dissolving  the  salt,  that  causes 
the  fall  of  temperature  necessary  to  establish  the  equilibrium.  If  this 


380  INORGANIC  CHEMISTRY.  [§  258. 

fall  is  to  be  considerable,  the  solubility  of  the  salt  must  of  course  be 
great.  In  this  case  also  the  cryohydric  point  is  the  lowest  tempera- 
ture that  can  be  reached  by  the  mixture. 

Chloride  of  lime  is  a  name  given  to  a  product  obtained  by 
saturating  slaked  lime  with  chlorine  at  ordinary  temperatures. 
Just  what  compound  is  formed  here  is  not  yet  definitely  known 
although  the  matter  has  been  frequently  investigated.  There  is, 

OC1 

however,  much  evidence  in  favor  of  the  formula  Ca  <  ™  ,  accord- 
ing to  which  it  is  a  mixed  salt  of  hydrochloric  and  hypochloious 
acids.  At  any  rate  this  is  more  probable  than  the  supposition 
that  chloride  of  lime  is  a  mixture  of  calcium  hypochlorite  and 
calcium  chloride  (§  56),  for  it  is  not  possible  to  extract  any  chlo- 
ride of  calcium  from  it  with  alcohol,  although  this  salt  is  very  solu- 
ble in  alcohol,  and  almost  all  the  chlorine  is  expelled  by  a  current 
of  carbon  dioxide. 

Chloride  of  lime  is  employed  in  large  quantities  for  bleaching 
and  disinfecting  (bleaching-powder) .  It  is  an  incoherent  white 
powder  with  the  odor  of  chlorine  (on  account  of  decomposition 
by  the  carbon  dioxide  of  the  air.)  When  treated  with  hydro- 
chloric or  other  acids  it  yields  chlorine: 

Ca°C1+2HCl   =CaCl2  +H2O+C12; 
+ H2S04 = CaS04 + H20  +  C12. 

A  solution  of  chloride  of  lime,  when  mixed  with  a  cobalt  salt  and 
warmed,  evolves  oxygen.  This  reaction  can  be  regarded  as  primarily  an 
oxidation  of  CoO  to  Co2O3,  the  latter  then  yielding  oxygen  with  chloride 
of  lime  and  forming  CoO  anew.  The  cobaltous  oxide  would  thus  act 
as  a  catalyzer. 

Calcium  fluoride,  CaF2,  occurs  in  nature  as  fluor  spar  or  fluorite, 
forming  cubes,  which  are  often  fluorescent.  It  is  insoluble  in 
water.  It  fuses  at  red-heat  and  is  frequently  employed  as  a  flux 
in  metallurgical  processes.  It  can  be  obtained  artificially  by  treat- 
ing a  solution  of  calcium  chloride  with  sodium  fluoride,  NaF. 

Calcium  sulphide,  CaS,  is  made  by  heating  gypsum  with  char- 
coal. On  treating  the  mass  with  water  calcium  hydrosulphide, 
Ca(SH)2,  is  formed,  whose  aqueous  solution  loses  hydrogen  sul- 


§258.]  SALTS  OF   CALCIUM.  381 

phide  on  boiling.  Calcium  sulphide  (like  the  sulphides  of  barium 
and  strontium)  has  the  property  of  emitting  light  in  the  dark 
after  it  has  been  exposed  to  sunlight,  but  seems  only  to  show 
this  phenomenon  when  it  contains  traces  of  other  elements, 
such  as  vanadium  or  bismuth.  A  boiled  mixture  of  lime-water 
and  sulphur  is  coming  into  extensive  use  as  an  insecticide  under 
the  name  of  "  lime-sulphur  solution." 

Calcium  sulphate,  CaSO4  •  2H2O,  occurs  in  nature  as  gypsum 
(§  256) .  It  is  only  slightly  soluble  in  water.  We  also  find  calcium 
sulphate  in  nature  as  anhydrite,  which  has  no  water  of  crystal- 
lization and  is  very  difficultly  soluble  in  water.  Gypsum  passes  over 
into  this  anhydrous  modification  on  being  ignited.  However,  the 
reverse  transformation,  recombination  with  water,  does  not  take 
place,  or  at  least  proceeds  very  slowly,  so  that  ignited  gypsum  is 
said  to  be  "  dead-burnt."  If  the  dehydration  is  carried  out  at  a 
lower  temperature,  an  anhydrous  gypsum  is  obtained  which  is 
comparatively  easily  soluble  in  water  ("soluble  anhydrite")  and 
absorbs  water  very  rapidly.  In  addition  to  these  varieties  there 
is  also  a  "half-hydrate",  2CaS04-H20.  This  is  the  chief  constit- 
uent of  "plaster  of  Paris."  On  being  stirred  with  water  it  takes 
up  the  latter  rapidly  and,  like  the  soluble  anhydrite,  forms  the 
dihydrate,  CaSC>4-2H20,  whereupon  the  mass  becomes  hard.  This 
is  the  basis  of  the  application  of  gypsum  in  the  manufacture  of 
casts,  etc. 

The  "setting"  depends  upon  the  relatively  high  solubility 
(about  1%)  of  this  half-hydrate,  on  account  of  which  it  forms  a 
solution  supersaturated  as  to  gypsum  (CaSO/i  •  2H20 ;  solubility 
about  0.2%)  and  gypsum  is  deposited.  Another  very  essential  fac- 
tor in  the  setting  is  the  filamentary  character  of  the  precipitated 
gypsum,  a  property  which  is  entirely  lacking  in  the  case  of  calcium 
hydroxide,  for  which  reason  slaked  lime  does  not  hold  together. 

The  credit  of  having  explained  the  conditions  governing  the  exist- 
ence of  the  above-mentioned  modifications  as  well  as  of  having  deter- 
mined the  positions  of  their  transition  points  is  due  to  VAN'T  HOFF. 
The  investigation  was  rendered  the  more  difficult  because  of  the  retard- 
ation phenomena  which  obscure  the  true  situation.  It  was  found  that 
the  half-hydrate  is  to  be  regarded  as  a  metastable  modification,  because, 
for  one  reason,  the  temperature  at  which  it  goes  over  into  soluble  anhy- 
drite is  lower  than  that  at  which  the  dihydrate  loses  all  its  water,  while 
in  general  the  loss  of  water  by  hydrates  proceeds  step  by  step  with  rising 


382  INORGANIC  CHEMISTRY.  [§§258- 

temperature.  Moreover  the  greater  solubility  of  the  half -hydrate,  as 
compared  with  that  of  the  dihydrate,  is  an  additional  reason.  There  is 
thus  the  same  relationship  here  as  between  the  metastable  crystals 
Na2SO4-7H2O  and  the  salts  Na2SO4-10H2O  and  Na2S04,  except  that  in 
this  latter  case  the  transformation  from  metastable  to  stable  modifi- 
cation takes  place  very  easily  on  touching  the  heptahydrate  with  a 
crystal  of  the  decahydrate,  while  the  half-hydrate  of  calcium  sulphate, 
even  in  contact  with  the  dihydrate,  retains  its  identity  indefinitely. 

Calcium  nitrate,  Ca(N03)2,  results  from  the  decay  of  nitrog- 
enous organic  substances  in  the  presence  of  lime.  It  crystallizes 
with  four  molecules  of  water.  The  anhydrous  salt  deliquesces  in 
the  air  and  dissolves  readily  in  alcohol.  It  is  converted  into  salt- 
petre by  potash  or  potassium  chloride  (§  229). 

Calcium  phosphates. — The  tertiary  salt,  Ca3(PO4)2,  is  insoluble 
in  water,  as  is  also  the  secondary  salt,  Ca2H2(P04)2,  The  primary 
salt,  CaH4(P04)2,  however,  is. readily  soluble;  it  is  employed  in 
large  quantities  as  an  artificial  fertilizer,  under  the  name  of  "  super- 
phosphate." 

This  superphosphate  is  manufactured  by  thoroughly  mixing  ground 
phosphorite  (or  bone  meal)  in  a  cast-iron  mixer  with  chamber  acid 
according  to  the  proportions  of  the  equation 

Ca3(P04)2  +  2H2SO4  =  CaH4(P04)2 + 2CaS04. 

The  mass,  which  is  at  first  semi-solid,  soon  becomes  solid,  since  the 
calcium  sulphate  that  is  formed  takes  up  the  water  contained  in  the 
chamber  acid  to  form  crystals. 

When  superphosphate  is  mixed  with  soil  the  primary  calcium  sulphate 
goes  into  solution  and,  since  every  soil  contains  lime,  it  is  forthwith 
reconverted  into  insoluble  secondary  or  tertiary  phosphate.  Appar- 
ently nothing  has  been  gained  toward  "making  the  phosphoric  acid 
soluble."  However  the  phosphate  is  now  diffused  widely  in  the  soil 
and  is  therefore  much  more  accessible  to  the  roots  of  the  plants  than  if 
the  soil  had  been  mixed  with  tertiary  phosphate  only. 

259.  Calcium  carbonate,  CaC03,  is  dimorphous,  occurring 
rhombohedral  as  calcite  and  rhombic  as  aragonite.  When  the 
solution  of  a  calcium  salt  is  treated  with  soda,  calcium  carbonate 
is  at  first  precipitated  in  an  amorphous,  very  voluminous  and  more 
soluble  form;  after  a  short  time,  however,  it  turns  to  a  finely 
crystalline  powder.  It  is  very  slightly  soluble  in  water,  but  more 


259.]  SALTS  OF  CALCIUM.  383 

extensively  so  in  water  containing  carbonic  acid,  since  the  acid 
calcium  carbonate  is  then  formed.     The  latter  decomposes  when 
•  the   solution  is   boiled,   carbon   dioxide   escaping   and   crystalline 
neutral  carbonate  being  deposited. 

Hardness  of  Water. — Almost  every  river-  or  spring-water  holds 
more  or  less  lime  in  solution.  The  lime  is  present  as  sulphate  or 
as  acid  carbonate.  Such  a  water  forms  but  little,  if  any,  lather 
with  soap;  the  fatty  acids  of  the  soap  form  white  insoluble  salts 
with  the  lime,  so  that  water  containing  much  lime  is  not  good 
for  washing.  Such  a  water  is  termed  hard  in  contrast  with  a 
water  that  is  free  or  nearly  free  from  lime,  which  is  called  soft. 
If  the  hardness  is  due  to  acid  carbonate  (also  called  "bicarbonate" 
of  lime),  it  disappears  on  protracted  boiling,  calcium  carbonate 
being  precipitated.  In  such  a  case  we  speak  of  temporary 
hardness.  In  metallic  boilers  and  similar  vessels  the  carbonate, 
of  lime  that  is  deposited  adheres  firmly  to  the  sides  ("  boiler- 
scale").  If  the  hardness  of  a  water  is  due  to  gypsum,  which  is 
only  partially  removed  by  boiling  (§  236),  it  is  spoken  of  as 
permanent  hardness. 

When  heated,  calcium  carbonate  breaks  up  into  lime  and 
carbon  dioxide.  We  have  here  a  case  of  complete  heterogeneous 
equilibrium  (§  71),  for  the  susbtances  are  CaO  and  C02  and  the 
phases  CaO,  CaCOs  and  C02.  This  is  confirmed  by  experiments, 
which  show  that  the  concentration  of  the  gaseous  phase  (the  dis- 
sociation tension)  at  a  definite  temperature  is  constant  and  there- 
fore independent  of  the  amount  of  each  phase.  Complete  decom- 
position into  lime  and  carbon  dioxide  can  only  occur,  therefore, 
when  the  gaseous  phase  is  removed  (as  in  lime-burning,  §  257) 
or  when  its  tension  is  kept  below  the  dissociation  tension.  On 
the  other  hand,  if  the  tension  of  the  carbon  dioxide  is  greater 
than  the  dissociation  tension,  'calcium  carbonate  cannot  decom- 
pose. Under  these  circumstances  it  is  possible  to  fuse  calcium 
carbonate;  on  solidification  it  assumes  a 
crystalline  structure  and  becomes  marble. 

In  the  adjoining  Fig.  pn    let  AB  repre- 
sent the  dissociation  curve  of  calcium  car- 
"bonate   in   a   coordinate   system   Pt.     Only 
along   this   curve   are   the   three   phases   in 


FIG  60  equilibrium    with    each    other;     under    any 

other  conditions  one  of  the  phases  disappears 


384  INORGANIC  CHEMISTRY.  [§§  259- 

and  we  enter  either  the  region  of  the  phases  CaO  +  C02  or  that  of 
CaC03+C02. 

GLASS. 

260.  Calcium  silicate  is  chiefly  important  because  it  is  a  con- 
stituent of  almost  all  sorts  of  glass. 

Glass  is  a  mixture  of  silicates  of  the  alkalies  with  calcium  silicate 
or  lead  silicate.  The  alkali  silicates  are  soluble  in  water,  amorphous 
and  easily  fusible.  The  calcium  silicates,  however,  are  insoluble, 
very  hard  to  fuse  and  frequently  crystallized.  By  fusing  both 
together  an.  insoluble  amorphous  transparent  mass  of  moderate 
fusibility  is  obtained,  which  is  glass.  It  is  prepared  by  fusing  a 
mixture  of  clean  sand,  lime  and  soda  in  refractory  crucibles. 

The  properties  of  glass  depend  primarily  on  the  quality  of  the 
materials  and  secondarily  on  the  proportions  used.  By  varying 
these  two  conditions  it  is  easy  to  obtain  grades  of  glass  varying 
widely  in  fusibility,  hardness,  lustre,  refractive  power,  etc.  There 
are  very  many  different  sorts  in  use.  Some  of  the  most  important 
are  the  following: 

Soda  glass  (window-glass)  is  a  soda-lime  silicate.  It  is 
readily  fusible  and  is  used  for  most  purposes  of  the  household. 

Potash  glass  (crown  glass,  Bohemian  glass)  consists  of  a 
silicate  of  potassium  and  calcium.  It  is  very  difficult  to  fuse  and 
is  therefore  extensively  used  for  chemical  purposes  (combustion 
tubes,  etc.). 

Leadglass  (flint  glass)  is  a  silicate  of  potassium  and  lead. 
It  is  softer,  more  easily  fusible  and  highly  refractive  and  takes  on 
a  beautiful  lustre  when  polished.  It  is  therefore  used  for  optical 
instruments  and  fancy  glassware  ("cut  glass"). 

Besides  the  substances  mentioned  many  others  are  used  in  glass 
factories  to  impart  particular  properties  to  the  glass.  The  addi- 
tion of  boric  acid  or  the  partial  replacement  of  lead  with  thallium 
gives  lead  glass  a  still  higher  refractive  index.  An  admixture  of 
alumina,  A1203,  prevents  or  hinders  chemical  utensils  of  glass  from 
becoming  brittle  and  allows  the  replacement  of  part  of  the  alkali 
by  lime.  Certain  metallic  oxides  form  colored  silicates  and  are 
therefore  mixed  in  with  the  furnace  charge  to  color  the  glass 
(cobalt,  blue;  chromium  or  copper,  green;  uranium,  yellow- 
green  fluorescent,  etc.).  The  addition  of  bone-ash,  Ca3(PO4)2, 


260.] 


GLASS. 


385 


or  tin  oxide  gives  a  milky-white  opaque  glass.  The  following 
table  shows  the  percentage  composition  of  various  kinds  of  glass, 
as  determined  by  analysis: 


S1O2 

K2O 

NaaO 

CaO 

PbO 

AhOs 
and  Fe2Os 

Window-glass  .... 

70 

15 

13 

2 

Bottle-glass  
Crown  glass  

64 
74 

2 
19 

4 

21 

7 



9 

Flint  glass  

55 

14 

31 

Plate  glass  

72 

17 

6 

5 

Water  has  in  general  very  little  effect  on  glass;  nevertheless  it 
attacks  it  somewhat.  Old  window-panes  have  a  peculiar  irides- 
cence, due  to  surface  weathering'.  As  it  is  very  important 
in  exact  analyses  to  know  how  much  glass  can  be  dissolved  from 
the  utensils,  careful  investigations  have  been  carried  out,  the 
results  indicating  the  following:  When  the  glass  is  new  a  rela- 
tively large  amount  goes  into  solution;  this  amount  'gradually 
decreases  in  the  course  of  a  few  weeks  to  a  minimum.  At  the 
first  the  alkali  in  particular  is  dissolved  from  the  surface  and  the 
resulting  solution  then  acts  as  a  solvent  for  the  silicic  acid.  To 
prepare  glass  vessels  so  that  they  are  almost  wholly  unaffected 
by  water  they  are  subjected  to  a  jet  of  steam  for  a  quarter  of  an 
hour  or  left  for  several  weeks  full  of  water,  the  water  being 
renewed  occasionally.  Thus  there  is  formed  on  the  surface  a  thin 
layer,  rich  in  silica  and  lime,  which  protects  the  inner  portion 
from  the  action  of  the  water. 

The  dissolving  action  of  water  on  the  alkali  of  glass  can  be  readily 
shown  by  agitating  finely  powdered  glass  in  water.  The  liquid  at  once 
turns  phenolphthalein  bright  red. 

Glass  is  a  typical  amorphous  substance.  Such  substances  are 
often  defined  as  liquids  with  a  very  high  internal  friction  and  the 
behavior  of  molten  glass  on  cooling  is  an  excellent  illustration 
of  this  definition.  At  high  temperatures  molten  glass  is  a  thin 
liquid ;  if  the  temperature  is  allowed  to  sink,  the  consistency  .of 
the  glass  becomes  tougher,  so  that  between  the  wholly  liquid 


386  INORGANIC   CHEMISTRY.  [§§260- 

and  the  wholly  "  solid  "  states,  there  is  a  continuous  series  of 
half-liquid  states.  As  it  is  thus  impossible  to  find  a  temperature 
limit  to  the  applicability  to  glass  of  the  laws  of  solutions,  e.g. 
the  law  of  diffusion,  it  seems  rational  to  consider  the  "  solid  " 
amorphous  state  as  liquid,  in  contradistinction  to  the  crystalline 
state,  which  latter  is  truly  solid,  having  very  different  properties 
from  liquids. 

Solid  solution.— This  term  was  introduced  by  VAN'T  HOFF 
to  apply  to  a  solid  homogeneous  mixture.  The  best  example 
is  to  be  found  in  mixed  crystals,  including  isomorphous  mixtures 
(§  210).  Thus,  for  instance,  when  a  molten  mixture  of  silver  and 
gold  solidifies,  the  components  do  not  separate,  but  solidify 
together  in  homogeneous  crystals  of  the  same  composition  as 
the  melt.  (See  Fig.  68,  III.)  The  term  "  solid  solution  "  is 
applied  to  this  and  somewhat  similar  solid  mixtures,  because  they 
exhibit  some  of  the  properties  of  liquid  solutions,  e.g.,  in  mis- 
cibility  relationships.  Glass  represents  an  amorphous  type  of 
solid  solutions,  of  which  the  constituent  silicates  are  the  integral 
components,  but,  as  intimated  in  the  preceding  paragraph,  many 
are  inclined  to  regard  the  amorphous  solid  solutions  as  pseudo- 
solid  solutions,  i.e.,  really  undercooled  liquid  solutions. 


II.  Strontium. 

261.  This  is  one  of  the  very  widely  diffused  elements.  CLARKE 
showed  that  in  most  of  the  rocks  containing  calcium  this  latter 
metal  is  accompanied  by  small  quantities  of  strontium  and 
barium.  The  principal  strontium  minerals  are  strontianite, 
SrCO3,  and  celestite,  SrS04.  Its  compounds  are  very  analogous 
to  those  of  calcium. 

The  metal  has  been  obtained  by  the  electrolysis  of  fused 
strontium  chloride.  Its  specific  gravity  is  2.5.  In  its  properties 
it  corresponds  to  calcium  throughout. 

Strontium  oxide,  SrO,  is  formed  on  igniting  the  hydroxide  or 
carbonate.  The  temperature  required  for  the  complete  dissocia- 
tion of  the  latter  is  higher  than  that  for  the  corresponding  calcium 
compound.  The  hydroxide,  Sr(OH)2-SH2O,  is  more  soluble  in 
water  than  calcium  hydroxide.  The  chloride,  SrCl2-6H2O,  is 
hygroscopic,  like  that  of  calcium.  It  is  soluble  in  alcohol  and 


262.]  BARIUM.  387 

can,  with  the  aid  of  the  latter,  be  easily  separated  from  barium 
chloride,  which  is  insoluble  in  alcohol.  Strontium  sulphate  is 
much  less  soluble  than  calcium  sulphate;  at  16.1°  1  part  SrS04 
dissolves  in  10070  parts  H2O  (CaSO4,  1  part  in  543  at  15.2°).  In  a 
mixture  of  alcohol  and  water  it  dissolves  to  an  extremely  small 
extent .  Strontium  nitrate,  Sr(NOs)2,  is  insoluble  in  alcohol; 
this  forms  the  basis  of  separating  it  from  calcium  nitrate,  which 
dissolves  in  alcohol. 

Strontium  salts  are  used  in  pyrotechnics  because  of  the  beauti- 
ful crimson  color  they  impart  to  a  flame. 

III.  Barium. 

262.  This  element  occurs  combined  as  barite,  or  heavy  spar, 
BaSC>4,  and  as  witherite,  BaCOs,  in  considerable  quantities.  In 
preparing  the  other  barium  salts  it  is  merely  necessary  to  dissolve 
the  latter  mineral  in  the  proper  acid.  Barite,  however,  must  first 
be  reduced  by  ignition  with  charcoal.  This  can  be  accomplished 
in  the  electric  furnace: 

(1)  4BaS04+4C    =BaS+3BaS04+4CO; 

(2)  3BuSO4+BaS=4BaO+4SO2. 

The  metal  is,  in  this  case  also,  obtained  by  the  electrolysis 
of  the  fused  chloride.  Another  method  is  to  heat  the  oxide 
with  magnesium.  Barium  decomposes  water  vigorously  even 
at  ordinary  temperatures.  Sp.  g.  =  3.75. 

Barium  oxide,  BaO,  is  obtained  by  igniting  the  nitrate  or 
hydroxide  at  a  high  temperature.  It  unites  very  readily  with 
water  to  form  the  hydroxide,  Ba(OH)2,  which  is  rather  soluble  in 
water  (yielding  baryta-water),  and  crystallizes  from  the  hot  solu- 
tion on  cooling  in  pretty  lamince,  which  contain  eight  molecules  of 
water. 

Barium  peroxide,  BaC>2,  forms  on  heating  the  oxide  in  a  cur- 
rent of  oxygen  or  air.  When  it  is  introduced  into  dilute  sulphuric 
acid,  barium  sulphate  is  precipitated  and  hydrogen  peroxide  left 
in  solution.  If  baryta-water  is  again  added,  the  hydrate  BaC>2  •  8H2O 
crystallizes  out. 

Barium  chloride,  BaCl2-2H20,  is  not  hygroscopic  like  the 
chlorides  of  strontium  and  calcium.  The  nitrate  crystallizes 
anhydrous. 


388 


INORGANIC  CHEMISTRY. 


Barium  sulphate,  BaS04,  is  characterized  by  an  exceedingly 
small  solubility  in  water  and  acids;  at  18.4°  1  part  dissolves  in 
429,700  parts  H2O.  It  is  used  as  a  filler  and  as  a  pigment  under 
the  name  of  "  permanent  white/'  or  blanc  fixe.  Barium  carbonate 
yields  carbon  dioxide  only  at  very  high  temperatures,  prolonged 
heating  at  1450°  being  required  for  complete  decomposition. 

SUMMARY  OF  THE  GROUP  OF  THE  ALKALINE  EARTHS. 

The  following  small  table  summarizes  the  physical  properties 
of  the  elements  of  this  group: 


Be 

Mg 

Ca 

Sr 

Ba 

Atomic  weight  

9.1 

24.32 

40.09 

87.62 

137.37 

Specific  gravity  
\tomic  volume  

1.64 
5.6 

1.75 

13.8 

1.58 
25.2 

2.5 
34.9 

3.75 
36.5 

Color            

white 

white  . 

white 

white 

white 

As  to  the  specific  gravity  we  observe  that  only  in  the  cases  of 
Ca,  Sr  and  Ba  is  a  steady  increase  noticeable. 

In  respect  to  the  chemical  properties,  it  has  already  been 
remarked  that  these  elements  act  only  as  bivalent;  all  compounds 
of  the  group  therefore  have  the  same  formula  type.  In  the  solu- 
bility of  the  sulphates  a  gradual  decrease  is  to  be  observed  with 
rising  atomic  weight. 

Just  as  in  the  first  group  three  elements  K,  Rb,  Cs,  exhibit  a 
particular  kinship,  so  here  calcium,  strontium  and  barium  are 
closely  related  in  their  properties,  while  the  two  other  members  of 
the  group  are  unlike  them  in  many  respects.  Beryllium  displays 
analogy  with  aluminium  in  certain  points  just  as  lithium  does 
with  magnesium. 


263-]  SPECTROSCOPY.  389 


SPECTROSCOPY. 

263.  If  the  light  from  an  ordinary  gas  flame  or  the  WELSBACH 
incandescent  light  is  broken  up  by  a  prism,  there  is  projected  a 
continuous  series  of  perfectly  blended  colors  from  red  through 
yellow,  green,  and  blue  to  violet.  This  phenomenon  is  called  a 
spectrum,  and  since  it  is  unbroken,  a  continuous  spectrum.  We 
have  previously  remarked  that  the  luminosity  of  a  gas-flame  is  due 
to  incandescent  solid  particles  of  carbon.  It  has  been  found  to  be  a 
general  rule  that  incandescent  solids  give  a  continuous  spectrum. 

With  incandescent  gases  it  is  different.  If,  for  instance, 
we  split  up  the  light  from  a  Bunsen  flame,  in  which  salts  of  sodium, 
calcium  or  other  metals  are  volatilized,  we  see  only  a  few  narrow 
bands  of  light  in  certain  places,  the  rest  of  the  spectrum  being 
dark.  This  is  termed  a  line  spectrum.  Every  element  has  its 
own  peculiar  spectrum  lines.  If  the  spectrum  of  the  incandescent 
vapors  of  a  mixture  of  elements  is  carefully  examined,  it  is  found 


to  contain  all  the  characteristic  lines  of  each  element.  Since  it 
is  only  necessary  to  volatilize  extremely  small  amounts  of  sub- 
stances in  order  to  show  their  lines,  it  is  readily  seen  how  important 
the  spectrum-analytical  methods  introduced  by  BUNSEN  and 
KIRCHHOFF  must  be. 

For  the  examination  of  spectra  a  number  of  instruments  have 
been  constructed,  varying  according  to  the  particular  object  in 
view.  For  chemical  analysis  the  apparatus  of  VOGEL  or  that  of 
JOHN  BROWNING  is  now  very  generally  used.  It  is  a  small  direct- 
vision  spectroscope  which  gives  a  very  bright  spectrum  and  has  a 
sufficient  dispersion.  At  the  end  B  (see  Fig.  61)  is  the  slit  which 
can  be  made  narrower  or  wider  by  turning  the  rim  D.  The  small 
mirror  m  serves  to  throw  light  through  the  hole  P  on  to  an  auxil- 
iary prism,  in  order  to  compare  the  spectrum  of  the  light  which  is 
to  be  analyzed  with  that  of  a  known  source.  At  the  left  end 


390 


INORGANIC  CHEMISTRY. 


[§§263- 


is  the  ocular  through  which  the  spectrum  is  seen.     For  further 
information  text-books  on  physics  should  be  consulted. 

In  order  to  examine  the  spectra  of  metals  it  is  necessary  to 
convert  the  latter  into  the  form  of  vapor  at  a  high  temperature. 
There  are  different  ways  of  doing  this.  One  is  to  introduce  salts 
of  the  metals  into  a  colorless  flame  by  means  of  a  platinum  wire. 
The  heat  dissociates  halogen  salts  and  in  the  case  of  oxysalts 
converts  them  into  oxides,  which  are  reduced  to  the  metallic 
condition  by  the  hot  gases  of  the  flame.  This  methcd  is  very 
satisfactory  for  some  elements,  e.g.  those  of  the  alkali  and  alkaline 
earth  groups,  when  there  is  plenty  of  material.  In  other  cases  a 
flame  spectrum  of  this  sort  is  not  so  good  as  a  spark  or  an  arc  spec- 
trum, for  with  the  latter  it  is  possible  to  detect  with  accuracy 
extremely  small  amounts  of  a  substance.  Other  advantages  of  the 
latter  spectra  are  their  greater  light  intensity,  the  greater  con- 
venience in  execution,  and  the  like.  Moreover,  at  the  high  tem- 
perature here  prevailing  most  elements  exhibit  spectra  which 
cannot  be  obtained  with  the  gas-flame. 

A  spark  spectrum  can  be  obtained  in  a  very  simple  manner, 
thus:  Into  the  bottom  of  a  little  glass  cup 
(n,  Fig.  62),  about  15  mm.  wide  is  fused  a  platinum 
wire,  which  ends  in  a  tube  g  containing  mercury 
and  is  thus  connected  with  the  negative  pole  of  an 
induction  coil;  it  is  incased  in  a  conical  capillary 
tube  x,  beyond  which  the  wire  projects  about  0.5  mm. 
At  the  opposite  end  is  the  positive  electrode  in  the 
form  of  a  platinum  wire,  which,  with  the  excep-  \J\J}ri 
tion.  of  the  short  end  d,  is  fused  into  a  glass  tube  ; 
the  latter  is  fitted  into  the  cork  a.  If  some  of 
the  salt  solution  is  poured  into  the  cup  about  half 
way  up  the  negative  electrode,  the  liquid  is  drawn 
up  to  the  end  of  x  by  capillarity  and  every  spark 
volatilizes  a  tiny  portion.  In  this  way  there  is 
no  loss  of  material  and  the  sparks  are  very  uni- 
form, so  that  the  observation  of  the  spectrum  can 
be  continued  at  length. 

For  the  study  of  the  spectra  of  substances  which  are 
gaseous  at  ordinary  temperatures  the  P  L  i)  c  K  E  R-H  i  T  T  o  R  p 
(GEISSLER)  tubes  are  used  (Fig.  63).  The  gases  are  sealed  up 
in  them  in  a  very  dilute  condition.  On  connecting  one  of  these 


FIG.  62. 


264.]  SPECTROSCOPY.  391 

with  the  poles  of  an  induction  coil,  the  whole  tube  is  illuminated 
most  intensely  in  the  narrow  portion.  This  part  is  placed  ver- 
tically in  front  of  the  slit  of  the  spectroscope. 

Some  substances  have  the  property  of  absorbing  certain  colors 
and  transmitting  others.  If  the  solution  of  such  a  substance  is 
placed  before  the  slit  of  a  spectroscope  and  the  light  of  a  contin- 
uous spectrum  allowed  to  pass  through  it,  dark  bands  or  lines  are 
observed  in  the  spectrum.  A  number  of  substances  have  very 
characteristic  absorption  spectra. 

264,  The  spectroscope  is  one  of  the  most  delicate  means  we 
have  of  detecting  many  substances.  This  is  readily  seen  on  con- 
sidering how  small  an  amount  of  the  substance  under  examina- 
tion is  volatilized  by  the  sparks.  We  arrive  at  numbers  like 


FIG.  63 


0.3X10"6  mg.  sodium,  for  instance,  as  the  least  amount  that 
can  be  detected.  It  has  thus  been  possible  to  discover  elements 
which  occur  only  in  company  with  large  amounts  of  others  and 
would  therefore  have  been  very  difficult  to  find  in  the  ordinary 
way.  BUNSEN  and  KIRCHHOFF  themselves  found  caBsium  and 
rubidium  in  this  way  in  Durkheim  mineral  water,  In  order  to 
obtain  these  elements  from  it  in  the  form  of  chlorides,  it  was 
necessary  to  evaporate  44,000  kg.  water,  which  yielded  16.5  g.  of 
a  mixture  of  the  chlorides.  Other  very  rare  elements  which  were 
discovered  by  spectrum  analysis  are  thallium,  indium,  gallium, 
ytterbium  and  scandium. 

The  spectra  of  the  elements  differ  greatly  in  appearance,  as 
may  be  seen  at  once  from  Table  II  (Frontispiece).  The  numbers 
indicate  wave  lengths  of  light  expressed  in  hundredth  microns 
(10~5  mm.).  Certain  metals,  such  as  sodium,  thallium  and 
indium,  exhibit  only  one  distinct  line  when  their  flame  spectra 
are  examined  with  a  spectroscope  like  the  one  described  above. 
If  a  sparking  current  or  an  electric  arc  is  employed  for  the  vola- 
tilization of  the  substance  and  the  spectroscope  is  one  giving 
strong  dispersion,  many  more  lines  become  visible.  It  is  further 
found  on  photographing  spectra  that  there  are  still  more  lines 
in  the  infra-red  and  ultra-violet  portions,  which  are  invisible 
to  the  eye.  Present-day  spectroscopic  studies  deal,  therefore, 
almost  exclusively  with  carefully  prepared  photographs. 


392  INORGANIC  CHEMISTRY.  [§§264- 

The  number  of  spectral  lines  increases  rapidly  as  we  proceed 
to  elements  of  the  higher  groups  of  the  periodic  system.  While 
lithium,  sodium  and  potassium  give  20,  35,  and  41  lines,  re- 
spectively, the  spectrum  of  barium  contains  163  lines  and  that 
of  iron  more  than  5000  lines. 

Among  these  lines  there  are  certain  ones  which,  in  virtue  of 
their  position  (color)  and  intensity,  are  specially  characteristic 
of  an  element,  like  that  of  the  yellow  line  in  the  case  of  sodium, 
the  green  line  of  thallium  and  the  blue  lines  of  indium.  For 
purposes  of  identification  of  such  elements  these  prominent  lines 
are  generally  observed  directly  in  the  spectral  apparatus. 

Nitrogen  is  an  example  of  a  substance  that  gives  a  band  spec- 
trum when  it  is  examined  in  the  manner  described  in  §  263. 

265.  The  position  of  the  spectrum  lines  was  formerly  indi- 
cated numerically  according  to  an  arbitrary  scale.  Now  it  is 
expressed  with  the  aid  of  the  wave  length/},  10 ~7  being  taken  as 
a  unit  and  the  unit  being  called  the  ANGSTROM  unit  after  the 
physicist  who  introduced  it.  The  wave  length  of  the  sodium 
DI  line  was  found  to  be  expressed  by  5896.16  such  units.  The 
visible  part  of  the  spectrum  comprises  the  wave  lengths  of  about 
7500-4000  A.U. 

'Thanks  to  the  researches  of  ROWLAND,  MICHELSON,  KAYSER 
and  RUNGE,  and  others,  the  wave  lengths  of  a  very  large  number 
of  spectrum  lines  have  been  determined  with  great  exactness, 
so  that  one  is  encouraged  to  attack  the  question  whether  in  the 
apparently  very  promiscuous  distribution  of  lines  in  the  spectra 
there  is  such  a  thing  as  order. 

BALMER  was  the  first  to  show  that  this  is  the  case  in  the 

779 

hydrogen  spectrum.     The   formula  X=A—  2 — -,    where   A    is    a 

constant  (3646.13)  expresses  the  wave  lengths  X  of  the  lines 
of  the  spectrum  of  the  element  with  very  close  approximation, 
provided  ra  is  substituted  by  consecutive  whole  numbers  begin- 
ning with  3. 

The  spectra  of  other  elements  have  been  examined  for  similar 
regularities,  chiefly  by  RYDBERG  and  by  KAYSER  and  RUNGE, 
and  it  has  been  found  that  the  regularities  are  in  all  of  the  cases 
more  complex  than  for  the  hydrogen  spectrum.  It  would  lead 
us  too  far  to  enter  upon  a  discussion  of  these  questions,  which 


265.)  SPECTROSCOPY.  393 

properly  belong  to  Physics,  but  a  few  of  the  interesting  results 
are  worth  mentioning  here. 

RYDBERG,  who  has  devoted  particular  attention  to  the  spectra 
of  the  alkalies,  introduced  into  his  formulae  the  reciprocal  of  the 
wave  length,  the  oscillation  frequency  n,  which  represents  the 
number  of  wave  lengths  per  centimeter.  In  the  spectra  of  the 
alkalies  he  found  three  series  of  lines,  whose  oscillation  fre- 
quencies can  be  expressed  by  the  formula 


In  this  formula  NQ  is  a  constant  having  the  same  value  for  all 
these  metals  and  all  the  series;  UQ  and  ,«,  however,  are  two 
constants  that  have  different  values  for  each  series  of  each 
metal.  For  m  we  substitute  again  consecutive  integers,  as  in 
BALMER'S  formula.  This  last  formula  is,  moreover,  a  special 
case  of  that  of  RYDBERG,  since  BALMER'S  formula  can  be  trans- 
formed into 

1  M/ 


A  m* 

(where  n0'  =  A  and  NQ'  =  4A),  into  which  RYDBERG 's  formula 
is  also  transformed  when  j«  =  0.  The  values  of  the  constants 
of  these  different  series  were  found  by  RYDBERG  to  have  still 
further  definite  relationships. 

The  formula  of  KAYSER  and  RUNGE  isT  =  A  +  Bm~2  +  Cm~4, 

A 

in  which  A,  B  and  C  are  constants  and  m  consecutive  whole  num- 
bers. It  represents  the  wave  lerjgth  of  the  lines  in  many  cases 
more  faithfully  than  does  the  formula  of  RYDBERG;  however, 
there  is  no  relation  between  the  constants  A,  B  and  C  of  the 
different  series. 

The  spectral  lines  of  the  alkalies  also  exhibit  the  peculiarity 
of  consisting  of  double  lines  (doublets)  or  triple  lines  (triplets), 
the  wave-length  differences  being  constant  for  each  series. 

Similar  series  of  lines,  whose  oscillation  frequency  can  be 
expressed  by  one  of  the  above  formula?,  are  found  in  the  spectra 
of  some  of  the  elements.  In  the  spectra  of  many  others,  how- 


394  INORGANIC  CHEMISTRY.  [§§  265- 

ever,  they  are  lacking.  In  their  place  we  find  in  the  spectra  of 
lead,  tin,  arsenic,  bismuth  and  others  a  constant  difference 
between  the  oscillation  frequencies  of  a  considerable  number 
of  their  lines. 

Such  investigations  as  these  are  prompted  by  the  notion  that 
a  knowledge  of  the  laws  which  govern  the  distribution  of  the 
spectral  lines  of  one  and  the  same  substance  on  the  one  hand, 
and  the  variation  in  the  distribution  from  substance  to  sub- 
stance on  the  other  hand,  would  throw  some  light  on  the  nature 
and  kinetic  condition  of  the  atoms. 

With  the  aid  of  spectroscopy  it  has  been  possible  to  deter- 
mine what  elements  are  present  in  the  heavenly  bodies.  When 
light  from  the  latter  is  passed  through  a  prism,  line  spectra  are 
obtained  and  these  lines  correspond  in  position  to  those  of  terres- 
trial elements.  The  composition  of  sunlight  has  been  especially 
the  object  of  a  most  extensive  study.  The  spectrum  of  that  body 
contains  numerous  black  lines,  known  as  FRAUNHOFER 
lines.  The  theory  of  this  phenomenon  is  explained  in  Optics. 

By  comparing  the  FRAUNHOFER  lines  with  the  spectra  of  ter- 
restrial substances  it  has  been  found  that  the  sun's  atmosphere 
contains  chiefly  Fe,  Na,  Mg,  Ca,  Cr,  Ni,  Ba,  Cu,  Zn  and  H  (the 
latter  in  enormous  quantity).  Moreover,  for  450  lines  of  the  iron 
spectrum  there  are  found  to  be  corresponding  dark  lines  in  the 
sun's  spectrum.  On  the  other  hand,  the  solar  spectrum  displays 
countless  lines  which  are  not  yet  identified  in  terrestrial  spectra. 

We  are  led  to  presume  that  many  of  the  elements  to  which  these  lines 
are  due  will  also  be  revealed  on  the  earth  by  more  careful  research,  especially 
when  we  consider  what  a  small  part  of  the  earth  is  known  (see  footnote, 
p.  8).  This  presumption  has  been  strongly  confirmed  by  the  discovery  of 
helium  (§  111).  The  principal  line  of  the  latter,  Ds~ so  termed  because  of 
its  proximity  to  the  double  D-line  (D^)  of  sodium— was  observed  in  the 
spectra  of  many  fixed  stars  as  well  as  in  that  of  the  sun  before  the  element 
itself  was  identified  on  the  earth.  Helium  was  thus  discovered  in  the  stars 
before  it  was  found  on  the  earth.  It  is  a  striking  fact  that  it  occurs  in 
exceedingly  large  quantities  in  the  fixed  stars  (according  to  spectrometric 
observations)  while  there  is  apparently  only  a  very  small  amount  of  it  on  the 
earth. 


266.  THE  UNITY  OF  MATTER.]  395 


THE   UNITY  OF    MATTER. 

266.  The  notion  that  all  substances  are  derived  from  a  single 
original  substance  and  that  the  variety  that  we  observe  in  the 
material  world  is  merely  the  result  of  a  difference  in  arrangement 
and  form  of  the  smallest  particles  has  long  been  prevalent.  Even 
the  old  Greek  philosophers  had  a  fondness  for  it.  However,  the 
rise  of  experimental  investigation  was  not  very  conducive  to  the 
idea.  BOYLE  (1626-1691)  introduced  the  concept  element  in  its 
present  form.  According  to  this  concept  all  substances  are  to  be 
regarded  as  elements,  which,  with  the  means  at  our  command, 
cannot  be  further  resolved  into  dissimilar  components.  The  sub- 
sequent development  of  chemistry  has  shown  that  the  number 
of  these  elements  is  rather  large. 

Notwithstanding  that  the  idea  of  a  primordial  substance  lacked 
substantiation  and  was  more  or  less  discredited,  it  was  by  no 
means  rejected,  for  we  have  been  expressly  reminded  again  and 
again  that  the  substances  which  chemistry  regards  as  the  simple 
substances  are  only  classed  as  elements  conditionally;  the  pos- 
sibility always  remains  that  a  so-called  element  may  be  found 
capable  of  division  into  dissimilar  components,  as  has  often  ac- 
tually been  the  case. 

Impossible  as  it  was  to  deny  the  existence  of  a  primordial 
substance,  the  researches  of  the  18th,  and  a  large  part  of  those 
of  the  19th,  century  brought  to  light  nothing  to  support  the  idea. 
Not  until  the  discovery  of  the  Periodic  System  did  the  question 
again  demand  serious  attention.  This  discovery  was  the  first  to 
supply  an  experimental  basis  for  the  assumption  of  a  primordial 
substance  (§  220).  The  striking  dependence  of  the  properties  of 
the  elements  on  the  periodic  functions  of  their  atomic  weights, 
which  finds  its  expression  in  this  system,  leads  of  itself  to  the 
thought  of  a  fundamental  substance,  of  which  the  simple  sub- 
stances called  elements  may  be  said  to  be  polymers,  incapable  of 
resolution  by  the  means  at  our  command. 

Another  argument  for  the  divisibility  of  the  elemental  atoms 
is  contributed  by  spectroscopy.  In  order  to  explain  the  line 
spectra  exhibited  by  many  elements,  we  assume  that  the  move- 
ments of  the  atoms  give  rise  to  light  vibrations  of  definite  wave 


396  INORGANIC  CHEMISTRY.  [§  266- 

length,  which  are  perceived  by  us  in  the  ppectral  lines.  However, 
since  the  spectrum  of  a  single  element  is  extremely  complicated, 
we  should  have  to  assume  that  the  atoms  engender  very  complex 
movements.  The  simplified  hypothesis  was  then  offered  that  it 
is  not  the  entire  atom  but  smaller  particles  of  which  the  atoms 
are  composed,  that  give  rise  by  their  vibrations  to  the  different? 
spectral  lines. 

The  physical  investigations  of  the  last  decade  have  furnished 
substantial  reasons  for  believing  that  the  chemical  atoms  are  not 
in  reality  the  ultimate  particles  of  matter  but  that  they  are  divi- 
sible into  particles  approximately  2000  times  smaller  than  the 
hydrogen  atom.  These  particles,  carrying  with  them,  as  they  do, 
very  strong  electrical  charges,  are  called  "  electrons."  We  cannot 
in  this  book  do  more  than  indicate  the  main  observations  arid 
inferences.  The  hypothesis  of  electrons  is  the  result  of  the  study 
of  cathode  rays  in  connection  with  the  below-mentioned  investiga- 
tions of  radio-active  elements.  Cathode  rays  are  generated  when 
the  discharges  of  an  induction  coil  are  sent  through  a  rarefied  gas. 
It  is  assumed  that  from  the  cathode  there  are  projected  particles 
with  strong*  negative  charges,  electrons,  which  are  propagated 
with  a  velocity  of  several  thousand  kilometers  per  second.  Ac- 
cordingly, the  cathode  rays  consist  of  a  stream  of  these  electrons. 
Measurements  of  the  mass  of  an  electron  have  shown  that  it  is 
about  iniSnr  °f  that  °f  a  hydrogen  atom.  The  electrons  are  the 
same,  whatever  gas  is  contained  in  the  apparatus  and  whatever 
electrodes  are  used,  so  that  we  are  evidently  dealing  with  a  decom- 
position of  the  atoms  into  their  ultimate  components. 

The  anions,  according  to  the  electron  theory,  consist  of  atoms 
and  one  or  more  electrons;  the  chlorine  ion,  for  example,  con- 
sists of  chlorine  and  an  electron,  which  latter  is  represented  by  6. 
The  cations  are  formed  from  the  atoms  by  the  release  of  one  or 
more  electrons.  We  may  thus  write: 

C1'    and    K-0=K>. 


The  ionization  of  potassium  chloride  in  water  can  be  represented 
by  the  equation: 

KClaq=  (K-0) 


Certain  physicists  presume  to  be  able  to  go  a  step  further. 
Some  very  remarkable  investigations  of  ROWLAND  have  shown 


267.]  RADIO-ACTIVE   ELEMENTS,  397 

that  a,  moving  electrically  charged  conductor  exerts  the  same 
effect  as  an  electric  current,  so  that  the  latter  may  be  regarded 
as^consisting  of  very  swiftly  propagated,  discrete  electrical  par- 
ticles. Electricity  would  thus  have  an  " atomistic"  structure. 
FARADAY'S  law  points  in  the  same  direction.  According  to  this 
law  the  charges  on  the  ions  are  either  equal  to  or  a  multiple  of 
the  charge  on  a  hydrogen  ion.  Fractions  of  the  charge  are  not 
found.  This  is  a  very  significant  fact.  Just  as  for  the  explana- 
tion of  the  analogous  laws  of  DALTON  we  have  assumed  that  the 
elements  consists  of  atoms,  so  we  cannot  avoid  inferring  from 
FARADAY'S  law  that  electricity  is  divided  into  discrete  elemental 
particles,  which  are  to  be  regarded  as  "atoms  of  electricity." 
Furthermore,  it  has  been  shown  that  induction  and  other  pheno- 
mena proceed  just  as  if  electricity  had  mass,  for  electricity  has  the 
same  properties  of  inertia  as  ponderable  matter.  The  conse- 
quence is  that  matter  is  identified  with  electricity  and  the  elec- 
trons are  no  longer  to  be  regarded  as  electrically  charged  mass  par- 
ticles but  as  electrical  charges  themselves,  without  a  material  body. 
These  hypotheses  would  not  only  unify  matter  but  would  also  dispel 
the  time-honored  notion  that  energy  and  matter  are  distinct. 

9 

RADIO-ACTIVE    ELEMENTS. 

267.  BECQUEREL  discovered  that  uraninium  emits  a  peculiar 
sort  of  rays  which  are  propagated  in  a  straight  line  and  act  o-n 
a  photographic  plate,  but  are  not  reflected,  refracted,  or  polar- 
ized. When  gases  are  traversed  by  them  the  gases  become 
electrical  conductors.  Now  when  uraninite  (or  pitchblende,  a 
uranium-bearing  mineral  of  very  complicated  composition)  was 
investigated  as  to  its  radiation  the  strange  fact  was  brought 
out  that  the  radiation  of  the  mineral  is  4.5  times  as  powerful 
as  that  of  its  constituent  metal,  uranium,  although  only  50% 
of  the  mineral  is  uranium.  Uraninite  must  therefore  contain 
one  or  more  substances  having  a  stronger  radiating  power  than 
uranium.  We  are  indebted  principally  to  the  gifted  couple, 
M.  and  MME.  CURIE,  for  the  discovery  that  the  emission  of  these 
special  rays,  or  the  radio-activity,  is  due  to  the  presence  of  very 
small  amounts  of  elements,  hitherto  unknown  and  of  very  sur- 
prising properties. 


398  INORGANIC  CHEMISTRY.  [§§  267- 

The  only  means  of  control  in  the  separation  of  these  elements 
from  the  other  compounds  in  uraninite  after  the  removal  of 
uranium  was  to  measure  the  radio-activity  of  the  product  ob- 
tained in  each  operation.  This  was  accomplished  by  measuring 
the  conductivity  of  a  layer  of  air  that  was  exposed  to  the  rays. 
Thus  after  numerous  chemical  operations  the  active  substance 
was  concentrated  more  and  more.  This  method  is  comparable 
to  that  employed  by  BUNSEN  and  KIRCHHOFF  in  isolating  rubid- 
ium and  caesium  from  the  Diirkheimer  mineral  water,  where  the 
spectroscope  (§  232)  indicated  the  progress  of  the  concentration 
of  these  elements.  However,  the  measurement  of  radio-activity 
is  many  thousand  times  more  sensitive  than  a  spectroscopic 
examination.  Were  it  not  for  this  fact,  the  discovery  of  the 
radio-active  elements  would  have  been  impossible,  because  they 
occur  in  such  extremely  small  quantities.  For  example,  2000  kg. 
uraninite  residues  from  Joachimsthal  yield  only  about  0.2  g. 
radium  chloride. 

Radium  is  the  best  known  of  these  elements.  It  is  the  only 
one  that  has  been  isolated  and  whose  compounds  have  been 
prepared  in  the  pure  state.  Its  spark  spectrum  has  three  very 
bright  lines  in  the  blue  and  violet  and  accordingly  the  Bunsen 
flame  color  is  carmine.  In  its  chemical  behavior  it  shows  close 
analogy  to  barium;  it  is  separated  from  the  latter  element  by 
fractional  crystallization  of  the  bromides,  radium  bromide 
being  more  difficultly  soluble  than  the  corresponding  barium  salt 
(this  is  true  for  all  the  respective  salts  of  the  two  elements). 
With  the  aid  of  the  spectroscope  it  can  be  determined  whether 
the  salt  is  entirely  free  from  barium  bromide. 

The  atomic  weight  of  the  radium  thus  purified  was  found  to 
be  226.4,  which  could  not  be  raised  by  further  fractional  crystal- 
lization. With  this  atomic  weight  radium  fits  exactly  into  the 
second  group  of  the  periodic  system.  All  radium  salts  are  lumi- 
nous and  excite  a  large  number  of  substances,  such  as  barium 
platinocyanide,  BaPt(CN)4,  uranyl  sulphate,  precious  stones,  and 
the  like,  to  powerful  fluorescence.  It  similarly  affects  the  dia- 
mond. Genuine  diamonds  can  thus  be  distinguished  from  imi- 
tations. The  radio-activity  of  the  pure  bromide  is  about  a  million 
times  that  of  uraninite. 

MME.  CURIE  and  DEBIERNE  succeeded  in  1910  in  isolating  the 


§  297.]  RADIO-ACTIVE  ELEMENTS.  399 

element  itself.  They  electrolyzed  a  solution,  using  a  mercury 
cathode,  and  obtained  a  radium  amalgam,  from  which  the  mer- 
cury was  distilled  off  in  a  current  of  hydrogen.  Radium  is  a 
white  metal,  melting  at  700°.  Even  as  low  as  this  temperature 
it  volatilizes  appreciably.  It  is  attacked  by  the  air  and  decom- 
poses water  vigorously. 

Besides  uranium  and  thorium  the  most  important  radio- 
active elements  are  polonium,  actinium,  ionium,  and  radio- 
thorium.  Polonium  is  precipitated  in  a  number  of  reactions 
with  bismuth;  by  hydrogen  sulphide,  as  well  as  when  the  basic 
salts  of  bismuth  are  precipitated  by  water;  stannous  chloride 
precipitates  it  in  the  same  way  as  mercury  and  tellurium.  It 
is  also  deposited  on  a  rod  of  silver  or  bismuth  when  one  of  these 
is  immersed  in  a  solution  containing  polonium.  The  radio- 
activity of  polonium  is  about  a  thousand-fold  as  great  as  that  of 
radium.  From  15  tons  of  pitchblende  MARCKWALD  could  only 
obtain  3  mg.  polonium  salt,  still  somewhat  impure;  so  that  polo- 
nium also  surpasses  radium  considerably  in  scarcity.  Actinium 
occurs  with  the  rare  earth  metals,  particularly  lanthanum,  and 
can  be  partially,  though  unsatisfactorily,  separated  from  them  by 
fractional  crystallization  of  the  manganese  double  nitrate.  For 
ionium  see  below;  for  radiothorium  see  under  thorium. 

The  rays  emitted  by  radium  preparations  are  of  three  sorts 
and  are  distinguished  as  a-  ft-  and  ^-rays.  Quantitatively  the 
first  are  predominant.  All  of  them  have  the  above-mentioned 
properties  in  common;  they  differ,  however,  in  their  penetrating 
power  and  in  their  behavior  in  the  magnetic  field.  The  a-rays 
are  not  very  penetrating  and  are  only  slightly  deflected  in  a 
strong  magnetic  field.  A  sheet  of  aluminium  foil  0.1  mm.  thick 
almost  entirely  stops  their  passage.  Moreover  they  are  com- 
pletely absorbed  by  a  layer  of  air  a  few  centimeters  in  thickness. 
The  /?-rays  are  strongly  deflected  in  a  magnetic  field  and  consist 
of  rays  of  various  but  greater  penetrating  power;  some  kinds 
of  /?-rays  can  even  pass  through  an  aluminium  plate  1  cm.  thick. 
The  f-rays  are  scarcely  deflected  at  all  and  go  through  obstruc- 
tions with  ease,  several  centimeters  of  lead  being  insufficient  to 
stop  them;  they  form  only  a  small  part  of  the  total  radiation. 

The  interesting  thing  is  that  these  rays  are  analogous  to  those 
generated  by  electric  discharges  in  highly  rarefied  gases.  The 


400 


INORGANIC    CHEMISTRY. 


266- 


/3-rays  are  to  be  regarded  as  cathode  rays  of  great  velocity. 
They  consist  of  negative  electrons  which  are  propagated  with 
very  great  velocity,  some  almost  with  the  velocity  of  light 
(300,000  km.  per  sec.).  From  the  deflection  which  they  undergo 
in  an  electrical  field  and  a  magnetic  field  of  known  intensity  their 
mass  is  calculated  to  be  (as  in  the  case  of  the  cathode  rays) 
about  2  -0V  <)  of  that  of  a  hydrogen  atom.  The  velocity  of  these 
electrons  can  also  be  calculated  from  the  same  data.  Their 
enormous  velocity  explains  the  great  penetrating  power  of  /?-rays. 
The  a-rays  resemble  a  sort  of  radiation  which  is  also  obtained 
by  discharging  electricity  in  a  rarefied  gas,  viz.,  the  canal  rays  of 


rrn 


FIG.  64. — EFFECT  OF  A  MAGNETIC  FIELD  ON  THE  a-,  /?-,  AND  J--RAYS. 

GOLDSTEIN.  They  behave  as  positively  charged  projectiles 
hurled  at  a  great  velocity  (about  -fa  that  of  light).  Their  mass 
is  about  equal  to  that  of  a  hydrogen  atom,  or  much  greater  than 
the  mass  of  the  projectiles  formed  by  the  /?-rays  and  the  cathode 
rays.  Their  greater  size  and  relatively  small  velocity  explains 
their  slight  penetrative  power.  In  the  light  of  the  more  recent 
investigations  they  appear  to  consist  of  helium  atoms  bearing 
two  positive  charges  each,  or,  more  specifically,  having  lost  two 
electrons.  The  ?--rays  are  analogous  to  the  X-  or  RoENTGEN-rays. 
These  proceed  from  a  metal  plate  which  is  placed  in  the  path  of 
cathode  rays;  they  do  not  consist  of  a  stream  of  electrically 
charged  particles,  but  are  regarded  as  a  form  of  wave  motion  of 
the  ether,  which  originates  when  electrons  are  projected  with 
great  velocity  against  a  solid  body. 

The  manner  of  detecting  the  various  sorts  of  rays  follows 


§  267.]  RADIO-ACTIVE  ELEMENTS.  401 

readily  from  the  above  description  of  their  properties.  Use  can 
be  made,  for  example,  of  their  dissimilar  penetrative  power. 
Their  separation  in  a  magnetic  field  is  diagrammed  in  Fig.  64. 
While  the  f-rays  suffer  no  deflection,  the  a-rays  are  deflected 
to  one  side,  the  #-rays  very  much  to  the  opposite  side. 

According  to  what  has  been  recited  in  the  preceding  sections, 
we  are  to  look  upon  these  forms  of  radiation  as  evidencing  a  spon- 
taneous decomposition  of  the  atoms  of  the  radio-active  elements. 
The  decomposition  is  accompanied  by  a  very  considerable  evolu- 
tion of  heat.  One  gram  of  radium  gives  off  about  118  g.-cal.  per 
hour;  for  this  reason  radium  salts  have  a  higher  temperature 
than  their  surroundings.  Even  cooling  with  liquid  hydrogen 
(—253°)  does  not  stop  this  evolution  of  heat.  The  magnitude 
of  the  heat  effect  is  more  apparent  upon  comparison  with  other 
caloric  effects  attending  chemical  reactions.  We  now  assume 
that  the  heat  evolution  in  the  decomposition  of  1  g.  radium  is 
about  109  g.-cal.  On  the  other  hand,  the  formation  of  1  g.  water 
from  its  elements  evolves  4X103  cal.,  so  that  the  first-mentioned 
process  gives  off  250,000  times  more  heat  than  the  second. 

The  spontaneous  decomposition  of  radio-active  substances 
is  accompanied  by  other  phenomena.  Every  substance 
that  is  brought  into  proximity  with  a  radium  salt  acquires  a 
temporary,  or  induced,  radio-activity,  i.e.,  it  emits  the  same  rays 
as  radium  itself.  This  induced  radio-activity  is  best  observed  on 
putting  a  radium  salt  in  an  enclosed  space.  The  enclosing  walls, 
as  well  as  all  bodies  within  the  space,  become  active.  It  is  not 
the  radium  rays  that  cause  this  effect,  for  a  radium  salt  in  a 
sealed  tube  emits  rays  without  exciting  any  radio-activity. 
RUTHERFORD  discovered  the  cause  of  this  phenomenon  by  the 
observation  that  there  is  a  constant  outflow  from  radio-active 
substances,  which  outflow  he  calls  emanation. 

Since  bodies  with  induced  radio-activity  give  out  rays  that 
are  identical  with  those  of  radium  itself,  these  rays  can  be 
regarded  as  transformation  products  of  the  emanation  of 
radium. 

Emanation  behaves  in  many  respects  as  a  gas;  it  diffuses  from 
one  vessel  into  another,  follows  the  law  of  BOYLE  in  its  compres- 
sion, can  be  condensed  by  cooling  with  liquid  air  and  volatilized 
.again  if  the  temperature  is  allowed  to  rise.  Neither  physical  nor 


402  INORGANIC  CHEMISTRY.  [§  267. 

chemical  agencies  are  able  to  alter  emanation.  It  is  indifferent 
to  temperature  variation  between  -180°  and  500°,  is  not  absorbed 
by  concentrated  acids  or  alkalies,  and  can  be  conducted  without 
change  over  hot  copper  oxide.  It  has  the  properties  of  a  gas  of 
the  argon  group.  Ramsay  has  lately  succeeded  in  preparing 
radium  emanation  in  somewhat  larger  quantities.  He  calls  this 
element  niton.  It  is  found  to  be  a  water-clear  liquid  with  a 
specific  gravity  of  about  5  and  a  boiling-point  of  —62°  under 
atmospheric  pressure.  In  glass  vessels  it  is  very  highly  fluores- 
cent. 

Emanation  and  induced  radio-activity  must  be  considered 
as  intermediate  stages  in  the  complete  disintegration  of  the 
radium  atom  into  the  above-mentioned  radiations.  How- 
ever, other  substances,  part  of  which  have  not  been  further 
studied,  are  formed  simultaneously.  One  of  them  is  pretty 
well  known,  viz.,  helium.  RAMSAY  and  SODDY  have  demon- 
strated that  helium  is  formed  in  the  spontaneous  decom- 
position of  radium  emanation.  The  maximum  quantity  of 
emanation  that  could  be  obtained  from  50  mg.  radium  bromide 
was  conducted  by  them  with  the  help  of  an  oxygen  current  into 
a  U-tube  cooled  by  liquid  air  and  the  U-tube  was  then  evacuated 
with  a  pump.  A  vacuum  tube  which  was  fused  on  to  the  U-tube 
showed  no  traces  of  helium  after  removal  of  the  liquid  air.  The 
spectrum  appeared  to  be  that  of  an  unknown  element — presumably 
emanation.  After  the  apparatus  stood  four  days  the  helium 
spectrum  appeared.  This  phenomenon  explains  the  mysterious 
persistent  occurrence  of  helium  in  radium-bearing  minerals. 

The  law  governing  the  rise  and  decay  of  radio-active  sub- 
stances is  the  same  as  that  for  a  unimolecular  reaction.  As  we 
have  seen  in  §  50,  the  velocity  S  of  such  a  reaction  can  be  repre- 
sented by  the  equation: 


if  we  let  C  stand  for  the  concentration  and  K  for  a  constant. 
With  the  aid  pf  higher  mathematics  this  equation  can  be  trans- 
formed into: 

C=Cne~Kt. 


§267.]  RADIO-ACTIVE  ELEMENTS.  403 

where  CQ  is  the  initial  concentration  and  e  the  base  of  natural 
logarithms.     The  logarithmic  form  of  the  equation  is  : 

l~=-Kt. 


This  same  equation  holds,  as  above  stated,  for  the  velocity  of 
decomposition  of  a  radio-active  substance;  in  that  case,  however, 
we  understand  by  C  the  intensity  of  radiation.  This  magnitude 
can  be  determined  electrometrically. 

If  a  radio-active  substance  changed  into  only  one  new  sub- 
stance, the  phenomenon  would  be  very  easy  to  represent  graph- 
ically; for  upon  plotting  the  time  on  the  abscissa  axis  and  the 
logarithm  of  the  activity  on  the  ordinate  axis  the  phenomenon 
would  be  represented  by  a  straight  line.  But  when  the  substance 
A  is  converted  into  another  active  substance  B,  and  this  again 
into  a  new  active  substance  and  so  on,  the  situation  becomes 
much  more  complicated.  A  graphic  representation  with  the 
same  co-ordinates  as  before  would  no  longer  yield  a  straight 
line,  but  a  rather  complicated  curve.  Nevertheless,  it  has  been 
found  possible  to  resolve  these  experimental  curves  and  to 
calculate  with  certainty  the  number  of  active  substances  which 
participate  in  the  transformations,  as  well  as  their  constant  K. 
This  is  not  the  only  method  of  ascertaining  the  number  and  kind 
of  the  intermediate  products.  We  can  often  distinguish  the 
individual  substances  involved,  by  a  study  of  the  kind  of  radiation 
given  off,  certain  of  the  substances  emitting  only  ct-rays,  others 
only  /?-rays,  and  still  others  a  mixture  of  all  three  rays; 
indeed  there  are  some  of  the  substances  which  emit  no  rays 
at  all. 

In  some  instances  these  active  substances  have  been  actually 
separated  by  physical  or  chemical  means.  Certain  of  the  sub- 
stances are  found  to  be  gaseous;  others  form  a  deposit  on 
solid  bodies.  The  gaseous  substances  can  be  condensed  by 
cooling. 

The  best  way  to  characterize  the  various  radio-active  sub- 
stances is  by  the  exponent  K  of  the  above  equation;  this  constant 
is  to  a  large  degree  independent  of  temperature  and  pressure; 
which  is  not  true  of  ordinary  reactions.  Frequently,  however, 


404  INORGANIC  CHEMISTRY.  [§267. 

use  is  made  of  another  magnitude,  related  to  K,  viz.,  the  period 
of  half  decay.  If  in  the  equation 

i., 

~K 

we  take 

<L  i 

C0    2' 

we  have  the  case  where  the  intensity  of  the  radiation  has  decreased 
to  just  half.  Solving  for  t,  we  have 

J±-l-2-f 
K~  K     t} 

as  the  length  of  time  necessary  to  reduce  the  intensity  to  half. 
It  is  this  magnitude  t  that  often  serves  as  a  characterizing  constant 
instead  of  K. 

As  a  result  of  such  observations  and  determinations  a  series 
of  transformations  of  radium  has  been  worked  out,  as  given  in 
the  table  on  page  405. 

Polonium,  as  we  see,  is  brought  into  relationship  with  radium. 
The  transformation  of  polonium  into  lead,  however,  is  far  from 
being  established. 

In  the  other  (parental)  direction  radium  is  related  to  uranium. 
Since  the  half-decay  period  of  radium  is  about  1760  years  and 
the  age  of  the  solid  earth-crust  is  counted  in  millions  of  years, 
the  radium  in  the  earth  would  long  since  have  disappeared,  if 
it  had  not  steadily  been  re-formed.  It  is  now  definitely  estab- 
lished that  uranium,  a  substance  much  slower  in  its  transforma- 
tions, is  the  parent  substance  of  radium.  One  of  the  most  signifi- 
cant evidences  of  this  is  that  in  the  various  uraniferous  and 
radium-bearing  ores  the  ratio  of  uranium  to  radium  is  very 
nearly  the  same.  In  order  to  prove  this,  BOLTWOOD  and  STRUTT 
determined  the  uranium  content  of  the  ores  by  ordinary  analytical 
methods  and  the  radium  content  by  collecting  the  emanation 
evolved  on  dissolving  the  ores  and  measuring  its  activity  with 
an  electrometer.  Since  the  decomposition  of  uranium  proceeds 
much  more  slowly  than  that  of  radium,  the  constant  relationship 
shows  that  the  radium  is  formed  from  the  uranium.  Indeed, 


§  267.] 


RADIO-ACTIVE  ELEMENTS. 


405 


Transformation 
Products. 

Physical  and  Chemical  Properties 

Time  of  Half 
Decay. 

Rays. 

Effective 
Limit  of 
Radiation  in 
the  Air. 

Uranium 

— 

4.6Xl09yrs 

a 



I 

Uranium  X 

— 

21.  5  days 

fcr 

— 

Ionium 

Produces  Ra 

^  104  yrs. 

«,/? 

2.87cm. 

I 

Radium 

At.  wt.,  226.5.     Characteris- 

I 

tic     spectrum  ;       produces 

1760  yrs. 

a 

3  .  5  cm. 

1 

helium  continuously 

(Niton) 

At.  wt.,  <  2.22.     Inert,  con- 

3. 86  days 

a. 

4.23  cm. 

i 

densable  gas 

Radium  A 

Deposit  of  induced  activity 

3  min. 

a 

4  .  83  cm. 

I 

Radium  B 

Same 

26.7  min. 

Slow 

— 

I 

• 

,5-rays 

Radium  d 

Same;     separation    by  elec- 

19.5  min. 

a,/?,  f 

a=7.06cm. 

I 

trolysis 

Radium  C2 

— 

1-2.  5  min. 

Ar 

— 

Radium  D 

One  of  the  constituents  of 

12  yrs. 

none 



radiolead.    Resembles  lead 

I 

closely  in  properties 

Radium  E^ 

Separable  from  Ra/)  by  elec- 

6.2 days 

/? 

— 

i 

trolysis 

Radium  E2 

— 

4.8  days 

/?  ' 

— 

Radium  F  or 

Sulphide  insol.  in  acids;  pre- 

140 days 

a. 

3.86cm. 

Polonium 

cipitated    with    basic    Bi- 

1 

salts;   pptd.  by  SnCl2 

Lead(?) 

_ 

_ 

~ 

SODDY  finally  succeeded  in  showing  that  solutions  of  most  care- 
fully purified  uranyl  nitrate  came  to  contain  radium  in  the 
course  of  three  years. 

Furthermore,  it  is  equally  well  established  that  there  are 
intermediate  products  between  uranium  and  radium;  in  other 
words,  that  uranium  is  not  the  direct  parent  of  radium.  One 
of  these  intermediate  products  is  ionium.  It  is  closely  allied 
chemically  to  thorium.  It  emits  a-  and  /5-rays,  of  which  the 


406  INORGANIC  CHEMISTRY.  [§§267- 

former  are  characterized  by  an  especially  feeble  penetrative 
power,  their  effective  limit  in  the  air  being  less  than  3  cm.  The 
percentage  relationship  between  ionium  and  radium  in  the 
different  ores  is  practically  constant.  The  most  interesting 
property  of  ionium  is  that  it  can  produce  radium. 

As  for  radium  and  uranium,  so  for  thorium  and  actinium,  series 
of  successive  decomposition  products  have  been  worked  out. 

Chemical  Effects  of  Radioactive  Substances. — Various  chemical 
reactions  are  brought  about  by  the  influence  of  radioactive 
substances.  Among  the  many  which  have  been  observed  we 
may  mention  the  conversion  of  oxygen  into  ozone  and  of  yellow 
phosphorus  into  red  phosphorus,  the  decomposition  of  iodic  acid 
and  the  dissociation  of  water  into  its  elements.  An  aqueous 
solution  of  a  radium  salt  is  constantly  giving  off  slight  amounts 
of  detonating-gas,  amounting  to  0.6  c.mm.  per  day  per  gram  of 
radium. 

Furthermore,  these  rays  have  the  property  of  developing  a 
strong  color  in  different  substances,  such  as  glass,  porcelain, 
and  the  alkali  salts;  the  same  color  can  be  produced  by  cathode 
rays.  The  skin  is  also  attacked  by  these  rays. 

Occurrence  of  Radioactive  Substances. — This  is  by  no  means 
limited  to  thorium-  and  uranium-bearing  minerals.  Indeed,  it 
has  been  shown  that  small  quantities  of  radioactive  substances 
occur  very  extensively  in  nature.  A  charged  electroscope 
gradually  loses  its  electrification  in  the  air, — a  phenomenon 
which  is  traceable  chiefly  to  the  ions  in  the  air.  Sea-water  also 
contains  slight  traces  of  radioactive  substances;  so  does  the 
earth  proper.  Many  springs  that  come  from  considerable  depths 
are  rich  in  radium  emanation;  in  this  respect  the  waters  of 
Gastein  (Austria)  and  Yellowstone  Park  are  particularly  noted. 

Radioactivity  of  Other  Elements. — The  radioactive  decay  of 
actinium  emanation  proceeds  5X101*6  times  faster  than  that  of 
uranium.  The  question  at  once  arises  whether  there  may  not 
be  other  substances  with  a  rate  of  decay  very  much  slower  than 
that  of  uranium.  We  could  then  conclude  that  there  is,  after 
all,  no  essential  difference  between  the  radioactive  and  the  other 
elements,  but  that  all  of  them  suffer  decay,  even  though  in  most 
cases  the  decay  is  so  slow  as  to  escape  observation.  To  be  sure, 
this  is  by  no  means  proven,  but  there  are  two  reasons  for  such 


268.]  ZINC.  407 

a  hypothesis.  In  the  first  place,  it  would  explain  why  so  many 
pairs  and  groups  of  elements  are  found  occurring  together  in 
nature:  niobium  and  tantalum,  for  instance;  selenium  and 
tellurium,  the  platinum  metals;  the  rare  earths.  The  two  last 
mentioned  are  groups  of  elements  whose  properties  are  quite  as 
closely  allied  as  those  of  the  decomposition  products  of  radio- 
active elements.  Secondly,  a  slight  radioactivity  has  been  detected 
in  potassium  and  rubidium. 

So  far  as  we  know,  the  radioactive  transformations  are  irre- 
versible. We  can  only  stand  by  and  look  on;  we  can  neither 
produce  nor  stop  them.  If  these  changes  should  prove  to  be 
a  general  property  of  matter,  it  would  mean  that  all  matter 
is  engaged  in  slow  decay.  Inasmuch  as  the  transformation  can- 
not be  made  to  go  backwards,  the  inference  could  be  drawn  that 
the  universe  was  constructed  out  of  a  primordial  substance  by 
some  sort  of  a  creative  act. 

ZINC. 

268.  The  most  important  zinc  minerals  are  calamine 
(H2Zn2Si05),sw^/isomfe  (ZnC03),  sphalerite,  or  blende  (ZnS),  and 
various  oxides.  The  principal  localities  are  Silesia,  England, 
Belgium,  Poland,  and  more  recently,  certain  districts  in  the  United 
States;  notably  southwestern  Missouri.  To  obtain  the  metal  the 
ores  are  roasted — the  gas  (S02)  from  the  sulphide  ores  is  con- 
verted into  sulphuric  acid — yielding  zinc  oxide.  In  the  older 
processes  this  is  mixed  with  coal  and  heated,  forming  carbon 
monoxide  and  zinc.  The  latter  distils  over  and  collects  in 
the  receiver  together  with  a  fine  gray  powder,  zinc  dust.  This 
"dust"  is  a  mixture  of  zinc  oxide  and  zinc  powder  and  is 
frequently  used  in  the  laboratory  as  a  vigorous  reducing- 
agent. 

The  metal  is  bluish-white  and  has  a  specific  gravity  of 
6.9-7.2.  At  ordinary  temperatures  it  is  brittle,  but  at  100-150° 
It  becomes  softer;  it  can  then  be  beaten  into  plates.  At  the  same 
time  the  specific  gravity  rises  to  7.2  and  the  metal  becomes  firmer. 
At  200°  it  again  becomes  brittle  and  can  be  easily  pulverized.  It 
melts  at  418°  and  boils  at  920°.  The  metallic  vapor  has  a  specific 
gravity  of  33.8  (H  =  l);  hence  its  molecular  weight  is  67.6. 
Since  the  atomic  weight,  as  deduced  from  DULONG  and  PETIT'S 


M)8 


INORGANIC  CHEMISTRY 


[§§  268- 


law,  is  65.4,  the  molecule  in  the  vaporous  state  can  contain  only 
one  atom.  The  same  is  true  of  the  related  metals  cadmium  and 
mercury.  Zinc  is  permanent  in  the  air,  since  it  becomes  firmly 
coated  with  a  protective  layer  of  oxide.  Zinc  dust  decomposes 
water.  When  heated  to  boiling  in  the  air  the  metal  burns  to  zinc 
oxide,'  producing  an  intensely  bright  light.  It  is  dissolved  very 
easily  by  hydrochloric  or  sulphuric  acid  with  the  evolution  of 
hydrogen;  it  is  an  interesting  fact,  however,  that  when  a  piece  of 
absolutely  pure  zinc  is  placed  in  either  of  these  acids  no  hydrogen 
is  generated.  If  this  piece  of  zinc  is  brought  in  contact  with  a 
platinum  wire,  effervescence  begins  at  once,  not  from  the  surface 


In 


FIG.  65. 


FIG.  66. 


of  the  zinc,  however,  but  from  that  of  the  wire,  and  zinc  goes  into 
solution.  Written  in  ions  the  process  is 

Zn  +  2H'  =Zn"  +  2H, 

and  its  explanation  is  just  the  same  as  that  given  in  §  203,  for  the 
formation  of  a  "lead  tree."  In  this  case  also  the  zinc  drives 
cations  into  the  solution  with  great  force,  itself  thus  assuming  a 
negative  charge,  with  which  hydrogen  ions  can  be  discharged. 
The  only  difference  seems  to  be  that  these  hydrogen  ions  discharge 
themselves  at  the  platinum  instead  of  at  the  zinc.  However,  this 
difference  is  not  real,  since  in  the  case  of  the  lead  tree  the  fresh 
portions  of  lead  are  deposited  on  the  outermost  parts  of  it.  The 
perfect  analogy  is  made  still  clearer  by  a  somewhat  modified  form 
of  the  experiment: 

When,  on  the  one  hand,  a  plate  of  amalgamated  zinc  and  one 
of  platinum  are  connected  by  a  metallic  wire  and  dipped  in  dilute 
sulphuric  acid  (Fig.  65)  hydrogen  is  evolved  from  the  platinum 
plate  and  when,  on  the  other  hand,  a  plate  of  amalgamated  zinc 
and  one  of  lead  are  similarly  connected  and  dipped  in  a  dilute  solu- 
tion of  lead  nitrate  (Fig.  66)  lead  crystals  are  deposited  not  on  the 


269.]  ZINC.  409 

zinc  but  on  the  lead.  In  both  cases  the  negative  charge  of  the  zinc 
goes  through  the  wire  to  the  other  plate,  on  which  the  ions  of 
hydrogen,  or  lead,  as  the  case  may  be,  can  discharge  themselves. 

Metallic  zinc — often  called  spelter  in  commerce — has  numerous 
iises^  For  instance,  zinc  plates  are  very  extensively  used  for  roof- 
ing. Iron  is  frequently  coated  with  zinc  to  prevent  rusting;  it  is 
then  known  as  galvanized  iron.  Further,  zinc  is  a  constituent  of 
many  alloys,  e.g.  brass  (§  242). 

269.  Zinc  oxide,  ZnO,  is  usually  prepared  by  igniting  the  basic 
carbonate.  On  being  heated  it  turns  yellow;  on  cooling,  the 
original  white  color  returns.  It  is  employed  as  a  pigment  under 
the  name  zinc  white,  or  Chinese  white. 

Zinc  hydroxide,  Zn(OH)2,  is  precipitated  by  alkalies  from  the 
solution  of  a  zinc  salt  as  a  white  gelatinous  mass,  soluble  in  the 
alkalies  as  well  as  ammonia;  however,  the  reason  is  different  in  the 
two  cases.  In  the  presence  of  alkalies  zinc  hydroxide  behaves  as 
a  weak  acid;  it  forms  Zn02"  anions  and  the  cations  2H*,  which 
yield  a  salt  Zn(OK)2  with  the  alkali  in  the  ordinary  way  (§  66). 
When  treated  with  ammonia,  however,  a  complex  zinc-ammonia 
ion  is  formed,  which  is  soluble. 

Zinc  chloride,  ZnCl2,  can  be  obtained  by  heating  zinc  in  a  cur- 
rent of  chlorine  or  by  dissolving  zinc  in  hydrochloric  acid  and 
evaporating  the  solution.  In  the  latter  case  some  oxychloride  is 
formed,  however.  Zinc  chloride  melts  on  heating  and  distils  at 
680°.  It  is  very  hygroscopic  and  is  often  used  for  splitting  off 
water  from  organic  compounds.  On  adding  zinc  oxide  to  a'  con- 
centrated zinc  chloride  solution  a  soft  mass  is  obtained,  which  soon 
becomes  hard  because  of  the  formation  of  the  basic  chloride 

OTT 

Zn<Xi    .    Ammonia  unites  with  zinc  chloride  to  form  various 
\j\. 

compounds. 

.  Zinc  sulphate,  ZnS04-7H20,  crystallizes  in  well-developed 
crystals,  which  are  isomorphous  with  the  analogous  compounds 
MgS04-7H20,  FeS04-7H20,  etc.  It  is  prepared  commercially  by 
carefully  roasting  zinc  blende. 

Zinc  sulphide,  ZnS,  is  completely  precipitated  by  hydrogen 
sulphide  from  solutions  of  its  salts  to  which  sodium  acetate  has 
been  added  to  neutralize  the  acid  set  free  from  the  zinc  salt.  In 
the  absence  of  sodium  acetate  it  is  still  partially  precipitated,  even 
from  solutions  of  the  neutral  salts  of  strong  acids. 


410  INORGANIC  CHEMISTRY.  [§§270- 


CADMIUM. 

270.  Cadmium  is  very  frequently  found  in  zinc  ores.     Being 
more  volatile  than  zinc,  it  distils  over  first  in  the  extraction  of 
such  ores.     It  is  obtained  pure  by  repeated  distillation  or  by  con- 
version into  the  sulphide,  which  is  insoluble  in  dilute  acids   and 
can  therefore  be  easily  separated  from  zinc  sulphide. 

Cadmium  is  a  white,  rather  soft  metal;  sp.  g.,  8.6;  m.-pt., 
322°;  b.-pt.,  778°.  It  is  unaffected  by  the  air  but  burns  on  heat- 
ing, forming  a  brown  cloud  of  oxide.  It  is  difficultly  soluble  in 
dilute  hydrochloric  and  sulphuric  acids,  but  readily  soluble  in 
dilute  nitric  acid.  The  cadmium  molecule  in  the  gaseous  state 
contains  only  one  atom. 

Cadimum  oxide,  CdO,  is  obtained  as  stated  above  and  also  by 
heating  the  carbonate  or  hydroxide.  It  is  an  amorphous  brown 
powder.  Cadmium  hydroxide,  Cd(OH)2,  is  insoluble  in  caustic 
potash  or  caustic  soda  but  soluble  in  ammonia,  on  account  of  the 
formation  of  a  complex  ion.  Cadmium  chloride,  CdCl2,  crystallizes 
with  two  molecules  of  water  and,  unlike  zinc  chloride,  can  be  dried 
without  decomposition.  Cadmium  sulphide,  CdS,  is  characterized 
by  a  bright  yellow  color  (it  is  used  as  a  pigment).  It  is  insoluble 
in  acids.  The  sulphate,  CdSC>4,  usually  crystallizes  out  of  its 
aqueous  solution  as  3CdS04-8H20.  There  is  also  a  salt 
CdS04-7H2O,  which  is  analogous  in  composition  to  the  sulphates 
of  magnesium,  zinc,  iron,  etc. 

MERCURY  (Quicksilver). 

271.  Mercury  is  the  only  metal  that  is  liquid  under  ordinary 
conditions.     It  occurs  in  nature  in  cinnabar,  HgS,  and  also  native. 
The  chief  localities  are  Almaden  in  Spain,  Idria  in  Illyria,  Mexico, 
Peru,  California,    China    and    Japan.     To  obtain  mercury  from 
cinnabar  the  latter  is  roasted  in  furnaces,   sulphur  dioxide  and 
mercury  being  formed.     The  mercury  vapor  is  condensed  either 
>  large  chambers  or  in  peculiarly  shaped  earthen   retorts,   or 
pipes,  called  aludels.     It  is  brought  on  the  market  in  75-lb.  iron 
flasks. 


271.]  MERCURY.  411 

The  commercial  product  is  not  pure,  containing  more  or  less  of  other 
metals  in  solution  (e.g.,  lead,  copper,  etc.).  Such  impurities  can  be  readily 
detected  by  the  fact  that  they  make  the  mercury  adhere  to  a  glass  vessel. 
A  suitable  process  of  purification  consists  in  letting  it  fall  in  fine  drops 
through  a  long  column  of  nitric  acid  (sp.  g.,  1.1),  as  in  Fig.  67.  The 
foreign  metals  are  thus  completely  dissolved,  while  almost  no  mercury  is 


FIG.  67. — PURIFICATION  OF  MERCURY. 

lost  by  solution,  because  these  foreign  metals  precipitate  mercury  from 
its  salt  solutions.  After  being  washed  with  water  the  metal  is  dried  and, 
if  absolute  purity  is  desired,  it  is  then  distilled  in  vacuo.  But  a  vacuum 
distillation  of  itself  is  insufficient,  for  some  lead  goes  over  with  it. 

Physical  Properties. — Mercury  solidifies  at  —39.4°  and  boils 
at  360°.  Even  at  ordinary  temperatures  it  is  somewhat  volatile, 
especially  under  reduced  pressure;  when  gold  leaf  is  suspended  in 
a  bottle  over  mercury,  for  instance,  it  eventually  becomes  white. 


412 


INORGANIC  CHEM IS TR  Y. 


271- 


The  metal  has  a  specific  gravity  of  13.595  at  0°.  The  vapor  density 
is  99.36  for  H  =  1 ;  hence  the  molecule  weighs  198.72.  This  number 
also  represents  the  atomic  weight,  as  has  been  found  from  molecular 
weight  determinations  of  many  volatile  mercury  compounds. 

Amalgams. — Many  metals  have  the  property  of  dissolving  in 
mercury  or  forming  compounds  with  it.  These  metal  solutions 
or  compounds  are  called  "  amalgams."  Besides  by  the  direct 
contact  of  the  two  metals  they  can  sometimes  also  be  obtained 
by  allowing  mercury  to  act  on  the  solutions  of  metal  salts,  e.g. 
silver  amalgam  can  thus  be  prepared.  Some  metals,  such  as  tin, 
dissolve  in  mercury  with  heat  absorption;  others  like  potassium 
and  sodium  with  great  heat  evolution  and  vigorous  action.  If  a 
great  excess  of  mercury  is  used,  the  amalgams  are  liquid,  other- 
wise solid.  Sodium  amalgam  is  exceedingly  firm  when  it  contains 
more  than  three  per  cent,  of  sodium. 

Opinions  have  been  divided  as  to  whether  amalgams,  and  metallic  alloys 
in  general,  are  mixtures  or  compounds.  In  order  to  solve  the  question  a 
study  has  been  made  of  the  freezing-point  curves  (cf .  §  237)  of  different  pairs 


FIG.  68. — TYPICAL  FREEZING-POINT  CURVES    OF  PAIRS  OF  METALS. 

of  metals.  Additional  information,  often  of  a  decisive  nature,  has  been 
afforded  by  the  microscopical  examination  of  the  surface  of  the  alloy  after 
having  been  etched  with  dilute  acid  and  polished,  the  individual  crystals 
being  generally  distinguishable  in  this  way.  As  a  result  of  these  investiga- 
tions it  has  been  found  that  the  freezing-point  curves  for  pairs  of  metals 
(binary  alloys)  may  be  of  three  different  types  (Fig.  68).  In  some  cases  there 
is  no  association  of  the  metals  and  the  curve  (I)  takes  the  same  form  as  the 
frtezing-point  curve  of  a  salt  solution,  having  a  eutectic  point.  In  others  a 


272]  MERCUROUS  COMPOUNDS.  413 

compound  may  be  formed  (II),  in  which  two  eutectic  points  are  observed, 
one  for  each  of  the  components  with  the  compound.  The  compound  has  the 
composition  corresponding  to  the  abscissa  for  the  maximum  of  the  curve 
between  the  eutectic  points,  the  freezing-point  of  the  pure  compound  being 
lowered  by  the  addition  of  one  of  the  components,  just  as  is  the  freezing-point 
of  a  pure  component  by  the  addition  of  the  other  component.  In  still  other 
typical  cases  the  components  may  form  mixed  crystals,  or  solid  solutions 
(§  260).  The  form  of  the  freezing-point  curve  (III)  is  very  instructive  in 
this  case.  When  mixed  crystals  can  be  formed  in  all  proportions  the  curve 
has  no  eutectic  point ;  every  liquid  phase  gives  crystals  of  a  definite  composi- 
tion corresponding  to  the  composition  of  the  liquid  phase.  The  freezing-point 
curve  is  uninterrupted  in  its  course. 

In  addition  to  the  above  three  types  we  may  have  various  combinations 
of  them.  Investigations  of  the  above  sort  have  shown  that  amalgams  of 
potassium  and  sodium  form  compounds,  such  as  Hg6Na  and  Hg2Na.  In  the 
amalgams  of  zinc  there  are  neither  compounds  nor  mixed  crystals.  Mixed 
crystals  are  formed  in  the  amalgams  of  tin,  lead,  and  cadmium. 

The  study  of  the  solid  products  of  the  cooling  of  molten  metallic  mixtures 
seems  at  first  somewhat  complicated,  because  we  may  have  not  only  the 
solids  that  result  from  the  slow  cooling  in  accordance  with  the  freezing-point 
curves,  but  we  may  also  have  solid  metallic  mixtures  formed  by  the  sudden 
chilling  of  a  hot  mixture.  This  possibility  is  rather  advantageous,  however, 
since  we  are  thus  enabled  to  fix  for  study  at  room  temperature  the  relation- 
ships prevailing  at  a  higher  temperature. 

Chemical  Properties. — At  ordinary  temperatures  the  metal  is 
not  affected  by  the  air;  at  higher  temperatures  it  takes  up  oxygen 
to  form  the  oxide  HgO,  which,  however,  splits  up  again  into  its 
elements  on  further  heating.  Dilute  hydrochloric  and  sulphuric 
acids  do  not  attack  it  at  ordinary  temperatures  and  dilute  nitric 
acid  acts  only  in  the  presence  of  nitrogen  dioxide  (see  §  127). 
Mercury  unites  instantaneously  with  the  halogens  and  sulphur. 

Mercury  forms  two  sets  of  salts,  ous  and  ic,  the  former  being 
derived  from  mercurous  oxide,  Hg2O,  and  the  latter  from  mercuric 
oxide,  HgO. 

Mercurous  Compounds. 

272.  Mercurous  oxide,  Hg2O,  is  dark  brown.  It  is  precipitated 
from  the  solution  of  a  mercurous  salt  by  caustic  soda.  It  decom- 
poses at  as  low  a  temperature  as  100°  or  in  the  light,  yielding 
mercuric  oxide,  HgO,  and  mercury. 

Mercurous  chloride,  Hg2Cl2,  calomel,  can  be  prepared  in  the 
wet  way  by  precipitating  a  dissolved  mercurous  compound  with  a 
chloride,  or  in  the  dry  way  by  subliming  a  mixture  of  mercuric 
chloride  and  mercury.  It  is  a  white  powder,  insoluble  in  water, 


414  INORGANIC  CHEMISTRY.  [§§272- 

but  turns  dark  in  the  light  on  account  of  the  separation  of  metal- 
lic mercury.  Ammonia  blackens  it  by  forming  a  mixture  of 
mercuric  ammonium  chloride,  NH2HgCl,  and  finely  divided  mer- 
cury: 

2HgCl + 2NH3  =  H2NHgCl + Hg  +  NH4C1. 

Calomel  is  frequently  used  as  a  medicament. 

The  vapor  density  of  calomel  has  been  found  to  be  117.6  (H=l), 
which  corresponds  to  the  molecular  formula  HgCl.  When  calomel 
evaporates,  however,  a  dissociation  into  HgCl2  and  Hg  occurs;  these 
products  unite  again  on  cooling,  but  they  can  be  previously  separated 
by  diffusion.  It  is  for  the  above  reason  that  the  vapor  density  was  found 
to  be  half  the  amount  calculated  for  Hg2Cl2;  hence  the  correct  formula 
of  calomel  is  Hg^C^. 

Here  also  BAKER  noted  the  influence  of  traces  of  water  (cf,  pp.  333, 
334).  According  to  his  investigations  thoroughly  dried  mercurous 
chloride  does  not  dissociate  on  volatilizing  and  gives  a  vapor  density 
which  corresponds  to  the  formula  Hg2Cl2. 

Mercurous  bromide  and  iodide  are  even  less  soluble  than  the 
chloride.  The  solubility  decreases,  as  in  the  case  of  silver,  with 
an  increase  in  the  atomic  weight  of  the  halogen. 

Mercurous  nitrate,  HgNOs,  is  formed  when  cold  dilute  nitric 
acid  acts  on  an  excess  of  mercury.  It  is  hydrolyzed  by  water, 

OTT 
a    yellow   basic   salt   Hg2<^rQ     being   deposited.     It   therefore 

dissolves  without  decomposition  only  in  dilute  nitric  acid.  The 
mercurous  ion  is  evidently  only  very  feebly  basic.  A  solution 
of  mercurous  nitrate  is  slowly  oxidized  by  the  oxygen  of  the  air  to 
the  mercuric  salt,  but  the  addition  of  a  little  mercury  reconverts 
it  into  the  lower  form. 

Mercuric  Compounds. 

273.  Mercuric  oxide,  HgO,  is  red  and  crystallized  when  pre- 
pared by  heating  mercury  or  mercury  nitrate,  but  yellow  and 
amorphous  when  precipitated  from  solutions  by  a  hydroxide  of 
potassium  or  sodium.  The  difference  between  these  lorms  seems 
to  be  due  only  to  a  difference  in  the  coarseness  of  their  grains. 
Mercuric  oxide  turns  black  on  heating  and  red  on  cooling. 

Mercuric  chloride,  HgCl2,  corrosive  sublimate,  is  manufactured 
on  a  large  scale  by  heating  a  mixture  of  common  salt  and  mercuric 


274.]  MERCURIC   COMPOUNDS.  415 

sulphate;  it  sublimes  over,  whence  its  name.  At  room  tempera- 
ture 1  part  HgCl2  dissolves  in  15  parts  H2O.  It  is  more  soluble 
in  alcohol.  The  acid  reaction  of  its  aqueous  solution  indicates 
hydrolytic  dissociation;  if  sodium  chloride  or  potassium  chloride 
is  added  to  the  liquid,  the  reaction  becomes  neutral  because  of  the 
formation  of  a  double  salt  HgCl2  •  KC1  •  H2O.  This  is  more  soluble 
in  water  than  sublimate  itself. 

Mercuric  iodide,  HgI2,  is  yellow  when  it  is  first  precipitated 
from  the  solution  of  a  mercuric  salt  by  potassium  iodide,  but  it 
soon  becomes  red.  If  this  modification  is  heated,  it  passes  over 
into  a  yellow  form  at  150°,  the  original  red  color  returning  on 
cooling,  however.  There  is  evidently  a  transition  point  here. 

A  similar  change  of  color  (red  to  brown)  is  observed  in  the  double 
salt,  Cu2I2  •  2HgI2,  even  at  a  rather  low  temperature.  On  cooling,  the 
red  color  promptly  reappears.  This  is  an  excellent  example  of  a  sub- 
stance whose  modifications  interchange  quickly  on  passing  the  tran- 
sition point.  Usually  the  transition  occurs  slowly. 

Mercuric  iodide  dissolves  readily  in  potassium  iodide  solution. 
NESSLER'S  solution, a  very  valuable  reagent  in  testing  for 
ammonia,  is  made  by  mixing  the  above  mercuric  iodide  solution 
with  caustic  potash.  It  should  be  noted,  however,  that  many 
organic  nitrogen  compounds  give  much  the  same  coloration  as 
ammonia  with  NESSLER'S  solution. 

Mercuric  cyanide,  Hg(CN)2,  is  obtained  by  boiling  Prussian 
blue  with  mercuric  oxide.  It  crystallizes  in 'fine  large  colorless 
crystals. 

274.  The  mercuric  halides,  in  contrast  to  the  other  salts  of  the  mer- 
curic ion,  are  only  slightly  ionized  in  aqueous  solution.  For  this  reason 
they  exhibit  some  peculiar  reactions.  On  mixing  a  mercuric  solution 
with  one  of  a  chloride,  for  instance,  considerable  heat  is  given  off  because, 
undissociated  HgCl2  molecules  are  formed,  while  the  mixture  of  solu- 
tions ordinarily  obeys  the  law  of  thermoneutrality  (§238,  2). — Again, 
if  mercuric  oxide  is  shaken  with  a  solution  of  chloride,  bromide,  or 
iodide  of  potassium,  the  liquid  becomes  strongly  alkaline  because  of 
the  liberation  of  potassium  hydroxide.  This  is  due  partly  to  the  slight 
ionization  of  the  mercury  halides  and  partly  to  the  combination  of 
the  latter  with  the  excess  of  alkali  halide  to  form  very  stable  alkali 


416  INORGANIC  CHEMISTRY.  [§§  274- 

mercuric  halides.  The  stability  of  these  complex  compounds 
increases  with  rising  atomic  weight  of  the  halogen. — The  same  cause 
explains  the  reverse  fact,  viz.,  that  the  halogen  compounds  of  mercury 
are  only  with  difficulty  decomposed  by  alkalies.  In  order  to  precipi- 
tate all  the  mercury  from  mercuric  chloride  a  large  excess  of  potassium 
hydroxide  must  be  employed;  mercuric  iodide  and  mercuric  cyanide 
cannot  be  decomposed  by  potassium  hydroxide  alone.  Mercuric  cyanide 
is  so  little  ionized  that  its  conductivity  can  hardly  be  measured;  hence 
it  does  not  give  any  of  the  ordinary  mercury  reactions,  except  the  forma- 
tion of  the  sulphide  (since  the  latter  is  so  very  insoluble) .  This  cyanide 
can  be  regarded  as  a  type  of  compounds  rendered  inactive  because  of 
non-ionization. 

This  low  ionization  also  explains  the  formation  of  mercuric  cyanide 
according  to  the  method  mentioned  above.  When  mercuric  ions  and 
cyanide  ions  are  brought  together,  even  in  extremely  dilute  solution, 
they  must  unite  to  form  Hg(CN)2  molecules.  The  union  of  these  ions 
necessitates  the  sending  of  more  of  them  into  solution  by  the  mercuric 
oxide  and  Prussian  blue,  and  so  the  process  goes  on  until  the  formation 
of  mercuric  cyanide  and  ferric  oxide,  Fe203,  is  complete. 

The  mercuric  halides  (especially  corrosive  sublimate)  are  very  strong 
antiseptics.  It  is  an  interesting  fact  that  in  this  respect  also,  they 
become  more  effective  as  their  ionization  increases.  The  chloride  is  a 
more  powerful  antiseptic  than  the  cyanide.  The  addition  of  metal 
chlorides  diminishes  the  ionization  of  sublimate  and  at  the  same  time 
reduces  its  disinfecting  ability. 

The  reason  why  the  mercuric  chloride  for  use  in  sublimate  tablets  is 
nevertheless  mixed  with  an  excess  of  common  salt  is  partly  that  the  sub- 
limate is  thus  dissolved  more  rapidly  and  also  because  such  solutions 
keep  longer  than  those  of  the  pure  sublimate,  especially  when  prepared 
with  well-water. 

Mercuric  nitrate,  Hg(NO3)2,  forms  basic  salts  very  readily;  on 
diluting  its  solution  in  nitric  acid  .with  water  there  is  deposited  a 
compound  Hg(N03)2-2HgO-H2O,  which  is  converted  into  pure 
mercuric  oxide  by  boiling  with  water.  This  shows  that  the  bivalent 
mercuric  ion,  also  is  very  feebly  basic. 

Mercuric  sulphate  is  not  soluble  in  water  but  is  converted  by 
the  latter  into  a  basic  salt.  In  the  presence  of  much  water  the 
yellow  compound,  HgSO4  •  2HgO,  is  formed.  With  the  sulphates  of 
the  alkalies  it  forms  double  salts,  e.g.  HgS04  •  K2SO4  -  6H2O,  which 
are  isomorphous  with  the  corresponding  double  salts  of  magnesium 
(§  255),  iron,  etc. 


275.] 


SUMMARY  OF  THE   GROUP. 


417 


Mercuric  sulphide,  HgS,  is  black  when  precipitated  from  solu- 
tion; on  being  heated  in  the  absence  of  air  it  sublimes  in  dark-red 
crystals,  which  are  similar  to  natural  cinnabar  and  are  used  as  a 
pigment  (vermilion). 

This  transformation  to  the  red  modification  also  occurs  when  black 
amorphous  mercuric  sulphide  is  left  in  contact  with  a  solution  of  alkali 
sulphide.  The  black  form  is  more  easily  soluble  than  the  red.  After 
some  time  red  dots  are  seen  in  the  black  mass  and  they  gradually  grow 
till  the  whole  mass  is  red. 


SUMMARY   OF   THE  GROUP, 

275.  Here  again  a  gradual  change  in  the  physical  properties  is 
to  be  seen  as  the  atomic  weight  rises.  The  following  small  table 
presents  a  few  of  the  constants: 


Be 

Mg 

Zn 

Cd 

Hg 

9  1 

24  32 

65  37 

112  40 

200  0 

1  64 

1  75 

6  9 

8  6 

13  6 

Mcltinsr-point      

>900° 

>651° 

418° 

322° 

—  39  4° 

Boiling-point  

>Zn 

920° 

778° 

360° 

1 

1 

1 

In  respect  to  chemical  properties  it  should  be  noted  that  all 
of  these  elements  are  bivalent,  except  that  mercury  can  be  con- 
sidered as  univalent  in  its  ous-compounds.  Their  sulphates  unite 
with  those  of  the  alkalies  to  form  double  salts  of  the  same  type, 
RS04-R2'SO4-6H2O  (R'  =  K,  Na,  NH4);  the  beryllium  double  salt 
alone  crystallizes  with  3H2O.  The  hydroxides  of  this  group  are 
soluble  in  ammonia  with  the  formation  of  complex  ions,  or  else  they 
yield  insoluble  metal-ammonia  compounds  (Hg). 

The  neutral  salts  have  a  tendency  to  go  over  into  basic  salts. 
This  is  especially  marked  in  mercury;  in  the  case  of  cadmium  it 
is,  strange  to  say,  very  weak. 

With  the  halogen  compounds  of  the  three  related  metals  Zn, 
Cd  and  Hg  the  electrolytic  dissociation  is  small;  it  decreases  as 
the  atomic  weight  of  the  metal  rises  and  is  very  slight  in  the  case 
of  mercury. 


418  INORGANIC  CHEMISTRY.  [§  276- 

ELECTROCHEMISTRY. 

276.  As  early  as  the  beginning  of  the  nineteenth  century,  when 
DAVY  isolated  the  alkali  metals  by  means  of  the  electric  current 
(§§  223  and  227),  there  was  known  to  be  an  intimate  relation 
between  electrical  and  chemical  phenomena.  BERZELIUS  even 
went  so  far  as  to  suppose  that  affinity  could  be  perfectly  explained 
by  assuming  that  the  atoms  are  electrically  charged  and  that  these 
charges  are  the  attractive  or  repellent  forces.  The  galvanic  element 
has  been  for  a  long  time  a  familiar  means  of  converting  chemical 
energy  into  electrical  energy.  However,  it  was  not  until  1889  that 
a  theoretical  explanation  of  the  connection  between  chemical  and 
electrical  phenomena  was  offered;  this  explanation  by  NERNST  is 
not  only  a  very  satisfactory  one,  but  it  also  affords  an  insight  into 
numerous  chemical  phenomena.  The  key  to  the  explanation  is 
the  concept  of  "electrolytic  solution  tension/7  which  has  already 
been  referred  to  in  a  few  instances  (§§  203  and  268). 

When  a  metal  comes  in  contact  with  the  aqueous  solution  of 
one  of  its  salts  a  difference  in  potential  develops  between  the  two. 
This  phenomenon  is  explained  by  NERNST  as  follows:  Just  as 
a  liquid  continues  to  evaporate  at  its  surface  until  the  pressure  of 
the  vapor  becomes  equal  to  the  vapor  tension  of  the  liquid,  so  a 
salt  must  continue  to  dissolve  in  water  (evaporation  and  solution 
being  analogous  processes)  until  the  osmotic  pressure  of  its  solu- 
tion balances  the  solution  tension  of  the  salt.  Now,  according  to 
NERNST,  every  metal  also  has  a  certain  tendency,  dependent  only 
on  its  chemical  nature,  to  force  its  atoms  into  solution  as  ions. 
This  force,  called  the  electrolytic  solution  tension,  comes  into  action 
when  the  metal  is  immersed  in  an  electrolyte  and  its  strength  is  the 
less,  the  more  cations  of  the  metal  are  already  in  the  solution.  The 
amount  of  cations  sent  into  the  solution  is  very  small,  as  experi- 
ment shows,— so  much  so  that  it  cannot  be  determined  by  the 
usual  chemical  means.  The  cause  of  this  is  not  that  the  solution 
tension  is  low, — on  the  contrary,  the  latter  is  often  very  large — 
but  that  an  equilibrium  is  very  soon  reached,  because,  notwith- 
standing the  low  concentration  of  the  ions,  they  carry  a  very 
high  electrical  charge  and  the  negatively  charged  metal  soon 
attracts  its  positive  ions  in  the  solution  with  such  force  that  just 
as  many  ions  are  precipitated  on  the  metal  as  are  sent  out  into  the 
solution.  If  P  represents  the  solution  tension  of  a  metal  and  p 


276.]  ELECTROCHEMISTRY.  419 

the  osmotic  pressure  of  the  cations  in  the  solution,  there  are  three 
possibilities  to  be  distinguished: 

(1)  P>p.    The   metal   then   behaves  like   a  salt  in   contact 
with  its  own  unsaturated  solution.     It  forces  cations  into  the  solu- 
tion of  the  electrolyte,  so  that  the  solution  becomes  positively 
charged  and  the  metal  has  to   take  on  a  negative  charge.    An 
equilibrium  is  soon  established.     However,  if  the  free  positive  and 
negative  electricities  acquired  by  the  electrolyte  and  the  metal  are 
conducted  away  by  a  connecting  wire  the  metal  will  again  send 
cations  into  the  solution,  and  this  action  will  continue  till  p  reaches 
the  value  of  P. 

(2)  P  —  p.     There  can  be  no  potential  difference. 

(3)  P<p.     In  this  case  the  metal  corresponds  to  a  salt  intro- 
duced into  its  supersaturated  solution.     Cations  are  now  deposited 
on  the  metal  and  charge  it  positively,  the  electrolyte  becoming 
negative.     Here  also  a  state  of  equilibrium  must  soon  arise  since 
the  negatively  charged  electrolyte  tends  in  turn  to  draw  the  posi- 
tive metal  ions  back  into  solution. 

The  relation  between  the  potential  difference  E  and  the  mag- 
nitudes P  and  p,  is  expressed  by  NERNST  with  the  equation: 

P  ,.,» 


in  which  R  is  the  gas  constant,  T  the  absolute  temperature,  n  the 
valence  of  the  metal  ions  and  loge  the  natural  logarithmic, 

If  for  R  in  this  equation  the  value  is  substituted  which  was 
calculated  in  §  34,  E  is  not  obtained  in  volts,  since  electrical  magni- 
tudes are  measured  in  other  units  than  those  there  employed.  In 
this  case  we  must  introduce  for  R  the  value  0.860  XlO~4.  If 
BRIGGS'  logarithms  are  to  be  used,  the  modulus  2.3025  must  also 
be  included.  The  value  of  E  in  volts  then  becomes: 

0.860XlO-4X2.3025Tlog  - 

F    __  P 
hi  =  ---  , 
n 

9jr  p 

or  #  =  io-4.—  log-. 

n      &  p 

From  this  equation  it  is  seen  that  E  increases  arithmetically  when 
p  decreases  geometrically.    For  example,  if  the  ionic  concentration 

IT 

is  reduced  to  one-tenth,  E  only  increases  by  —  X  10~4  volt.    It  is 

thus  seen  that  the  potential  difference  is  not  much  affected  by 


420  INORGANIC  CHEMISTRY.  [§276- 

changes  in  the  concentration  of  the  electrolyte,  even  though  they 
be  quite  large. 

On  bringing  together  two  different  metals  and  their  salt  solu- 
tions an  element,  or  cell,  of  the  DANIELL  type  is  obtained  (copper 
in  copper  sulphate  and  zinc  in  zinc  sulphate,  the  two  pairs  separated 
by  a  porous  partition).  The  electromotive  force  of  such  a  cell  is 
found  from  the  difference  of  the  two  values  of  E,  i.e. 


when  both  metals  have  the  valence  n. 

In  such  a  cell  with  closed  circuit  there  are  differences  of  potential  not 
only  between  metal  and  solution  but  also  between  the  two  liquids  and 
between  the  two  metals.  Experience  has  shown,  however,  that  both 
of  the  latter  are  very  small  in  comparison  to  the  former,  so  that  they 
may  be  disregarded. 

Leaving  the  solution  tensions  PI  and.P2  out  of  consideration, 
E  therefore  depends  on  the  values  of  the  osmotic  pressure  pi  and 
p2  of  the  metal-ions.  If  p2  can  be  made  extremely  small,  so  that 

P  P 

loge  —  <  loge  —  ,  E  becomes  negative,  i.e.  the   current  must  alter 

Pi  P2 

its  direction.     This  can  be  demonstrated  as  follows: 

In  a  DANIELL  cell,  in  which  the  osmotic  pressure  of  the  zinc 
ions  (pi)  is  seldom  very  different  from  that  of  the  copper  ions 
(p2),  the  current  goes  from  the  copper  through  the  connecting 
wire  to  the  zinc,  for  the  solution  tension  (Pi)  of  the  zinc  is  much 
larger  than  that  (P2)  of  the  copper  (see  below).  Now  the  con- 
centration of  the  copper  ions  can  be  made  several  powers  of  ten 
smaller  by  adding  potassium  cyanide  to  the  copper  sulphate  solu- 
tion, for  by  this  means  the  very  slightly  ionized  complex  (Cu2Cy4)" 
is  formed  (§  243).  This  addition  actually  reverses  the  direction  of 
the  current.  Neither  the  precipitation  of  the  copper  by  potassium 
hydroxide  nor  the  precipitation  by  ammonium  sulphide  reduces  the 
concentration  of  the  copper  ions  enough  to  produce  this  effect. 
Since  equation  (2)  can  also  be  written 


and,  when  p2=pi,  the  last  expression  becomes  zero,  it  is  apparent 
that  the  electromotive  force  of  a  DANIELL  cell  is  mainly  determined 


276.] 


EL  EC  T  ROC  HEM  IS  TR  Y. 


421 


by  the  ratio  of  the  solution  tensions  of  the  metals.  A  galvanic 
cell  can  be  regarded  as  a  machine  driven  by  the  electrolytic  solution 
tensions  of  the  metals. 

The  introduction  of  this  conception  of  solution  tension  and  the 
ideas  connected  with  it  has  led  to  an  altogether  clearer  understand- 
ing of  the  chemical  processes  of  galvanic  cells,  as  well  as  of  the 
way  in  which  the  current  is  generated  in  them. 

Galvanic  cells  may  be  divided  into  two  classes,  reversible  and 
non-reversible.  The  DANIELL  belongs  to  the  first  class.  It  pro- 
duces a  current  because  the  solution  tension  of  the  zinc  exceeds 
that  of  the  copper.  The  zinc 
sends  its  positi  vely  charged  ions 
into  the  sulphate  solution  and 
itself  becomes  negative.  On  the 
other  hand,  the  copper  ions,  on 
passing  over  into  atoms  and  pre- 
cipitating themselves  on  the 
copper  plate, transfer  their  posi- 
tive charges  to  the  latter,  which 
thus  becomes  the  positive  pole. 
Chemically  the  process  amounts 
to  the  simultaneous  solution  of 
zincandlprecipitationof  copper: 

CuSO4  +  Zn  =  Cu  +  ZnS04 
or,  in  ions: 


FIG.  69.— ACCUMULATOR. 


If  .a  current  is  sent  through 
the  DANIELL  cell  in  the  opposite 
direction,  ions  will  enter  into 
solution  at  the  copper  plate  be- 
cause the  latter  acquires  a  positive  charge,  and  the  zinc  ions  will 
be  forced  to  deposit  themselves  on  the  zinc,  for  the  reverse  current 
charges  the  zinc  negatively  so  that  it  attracts  the  zinc  ions.  It  is 
therefore  possible  by  passing  a  reverse  current  through  the  cell  to 
restore  it  to  its  original  condition — hence  the  term  "  reversible." 

One  of  the  most  important  styles  of  reversible  batteries  is  the 
accumulator,  or  storage-battery  (Fig.  69).  This  consists  of  a 
glass  jar  in  which  lead  plates  are  suspended  so  that  they  dip 
into  dilute  sulphuric  acid.  These  plates  are  coated  alternately 


422  INORGANIC  CHEMISTRY.  [§  276- 

with  lead  peroxide  (positive)  and  lead  sulphate  (negative). 
The  positive  plates  are  all  connected  with  each  other,  as  are 
also  the  negative  ones.  (From  a  large  number  of  such  cells 
a  battery  is  constructed  by  connecting  the  positive  pole  of 
each  cell  with  the  negative  pole  of  the  adjoining  one.)  If 
a  current  is  passed  through  the  system  so  that  it  enters  at  the 
lead  peroxide  plate  and  goes  through  the  sulphuric  acid  to  the 
other  plate,  lead  peroxide  collects  on  the  positive  plate,  while  on 
the  other,  the  cathode  plate,  the  lead  sulphate  is  converted  into 
spongy  lead.  By  this  process  the  accumulator  is  charged.  There- 
upon, if  the  poles  are  connected  (by  a  wire),  the  opposite  process 
goes  on;  the  lead  peroxide  is  reduced  at  the  one  plate  and  the 
spongy  lead  is  converted  into  lead  sulphate  at  the  other.  During 
the  discharge  the  peroxide  plate  is  again  positive,  the  lead  plate 
negative.  The  chemical  process  in  the  accumulator  cell  is  there- 
fore expressed  by 


The  generation  of  the  current  has  been  explained  in  various 
ways.  One  is  as  follows:  The  lead  peroxide  on  the  anode  plate  has 
a  certain  solution  tension,  and  hence  goes  into  solution  as  nega- 
tively charged  Pb(V'  ions.  Thereby  it  of  course  imparts  to  the 
plate  istelf  a  numerically  equivalent  positive  charge.  These 
bivalent  PbCV'  ions  encounter  positively  charged  Pb"  ions  at  the 
cathode  plate,  which  are  being  sent  by  it  into  the  solution;  the 
cathode  plate  charges  itself  negatively  at  the  same  time.  The  two 
sorts  of  ions  now  combine  to  form  electrically  neutral  PbO  mole- 
cules, which  yield  lead  sulphate  with  the  sulphuric  acid  present; 

Pb02"+Pb"  =2PbO;  2PbO  +  2H2SO4=2PbS04+2H2O. 

Among  the  non-reversible  cells  are  the  BUNSEN  and 
the  LECLANCHE.  .A  reverse  current  does  not  restore  these  to 
their  original  condition  and  their  electromotive  forces  E  cannot  be 
calculated  by  the  above  formula;  nevertheless  the  general  prin- 
ciples of  the  pressure  theory  can  be  applied  to  explain  the  produc- 
tion of  the  galvanic  current  in  these  cells. 

The  arrangement  of  the  BUNSEN  cell—  an  amalgamated  zinc 
plate  dipped  in  sulphuric  acid  and  a  carbon  cylinder  in  nitric  or 
chromic  acid—  is  well  known.  From  an  electrochemical  stand- 


277.] 


ELECTROCHEMISTRY. 


423 


point  the  generation  of  hydrogen  from  zinc  and  sulphuric  acid 
amounts  to  a  transfer  of  the  charges  of  the  hydrogen  ions  of  the 
dilute  acid  to  the  zinc  atoms  and  an  escape  of  hydrogen  in  the 
form  of  discharged  molecules.  In  the  BUNSEN  cell,  however, 
most  of  the  hydrogen  ions  find  an  opportunity  to  give  up  their 
positive  charges  to  the  carbon  cylinder  and  exercise  a  reducing 
action  on  the  nitric  or  chromic  acid.  On  the  other  hand  the  zinc 
plate  sends  positively  charged  zinc  ions  into  the  solution  to  the 
same  extent  as  hydrogen  ions  disappear,  the  zinc  plate  itself 
acquiring  a  negative  charge. 

The  LECLANCHE  cell  consists  of  a  zinc  bar  in  concentrated 
ammonium  chloride  solution  and  a  porous  earthenware  cylinder 
immersed  in  the  same  solution  and  containing  some  manganese 
peroxide  and  a  stick  of  carbon  for  conducting  off  the  current. 
Here  again  the  zinc  goes  into  solution: 

Zn + 2NH4C1  =  ZnCl2  •  2NH3  +  H2. 

The  hydrogen  ions  discharge  themselves  at  the  carbon  and  reduce 
the  peroxide.  In  this  case  also  the  carbon  is  the  positive,  the  zinc 
the  negative,  pole. 

277.  Just  as  in  galvanic  cells  chemical  energy  is  transformed 
into  electrical  energy,  so  reactions  between  ions  in  general  can 
produce  an  electric  current  if  the  conditions  are  suitable.  A 

few   examples    of    this   may  be 
cited. 

For  these  experiments  a  cell 
devised  by  LUPKE  is  very  satisfac- 
tory (Fig.  70).  It  consists  of 
two  glass  vessels  Zi  and  Z2,  to  the 
bottoms  of  which  the  platinum 
electrodes  ki  and  k2  are  attached. 
The  vessels  are  connected  by 
means  of  the  wide  siphon  H. 
The  wires  A  and  K  lead  to  a  gal- 
vanoscope.  To  show  that  elec- 
trical energy  can  be  obtained  by 
the  oxidation  of  the  stannous  to 
FIG.  70.-LtiPKE  CELL.  the  stannic  chloride  an  acidu- 

lated stannous  chloride  solution  (11.2  : 100)  is  introduced  into  Zi 


424  INORGANIC   CHEMISTRY.  [§  277- 

and  an  acidulated  normal  sodium  chloride  solution  into  Z2;  the 
siphon  also  is  filled  with  the  latter  solution.  As  soon  as  a  few  drops 
of  chlorine-water  or  a  solution  of  auric  or  mercuric  chloride  are 
allowed  to  fall  from  a  pipette  upon  the  electrode  (/c2)  in  the  salt 
solution,  the  galvanoscope  indicates  a  current  in  the  wire  circuit 
from  K  to  A.  Now,  in  order  that  the  bivalent  ion  Sn"  may  become 
quadrivalent  (Sn"")  it  must  acquire  two  more  positive  charges  and 
this  requires  the  addition  of  two  chlorine  ions.  These  are  at  once 
supplied  by  the  mercuric  or  auric  chloride.  The  metallic  ions 
(Hg"  or  Au"*)  are  deposited  on  k2  and  impart  to  the  latter  a  posi- 
tive charge,  which,  if  conducted  by  means  of  the  wire  circuit 
K,  is  at  the  disposal  of  the  Sn"  ions.  If  free  chlorine  is  added,  it 
splits  up  into  ions,  as  a  result  of  which  positive  electricity  is  imparted 
to  k2  and  this  flows  through  the  wire  circuit  back  to  ki  and  raises 
the  potential  of  the  Sn0'  ions. 

The  fact  that  electrical  energy  can  be  obtained  by  the  neutral- 
ization of  sulphuric  acid  is  capable  of  demonstration  with  the 
same  apparatus.  To  this  end  a  J-normal  sulphuric  acid  is  intro- 
duced into  Z2  and  a  ^-normal  potassium  sulphate  solution  into  Z\ 
and  the  siphon.  If  a  large  piece  of  palladium  foil  (about  4  sq.  cm.) 
that  has  been  saturated  electro lytically  with  hydrogen  is  placed 
on  the  platinum  disk  of  the  electrode  k\  and  touched  for  a  few 
moments  with  a  stick  of  caustic  potash,  bubbles  of  hydrogen 
will  rise  from  the  platinum  plate  of  the  other  electrode  k2  and 
the  needle  of  the  galvanoscope  will  indicate  the  passage  of  a  power- 
ful current  outward  from  k2.  The  hydrogen  of  the  palladium 
foil  sends  positive  ions  into  the  solution,  which,  however,  forth- 
with unite  with  the  OH-ions  of  the  potassium  hydroxide  to  form 
neutral  water.  By  the  emission  of  these  positive  ions  ki  acquires 
a  negative  potential,  which  flows  out  through  the  external  circuit 
to  k2.  The  hydrogen  ions  of  the  sulphuric  acid  surrounding  this 
electrode  are  thus  afforded  an  opportunity  of  discharging  them- 
selves against  this  negative  charge  so  that  hydrogen  is  given  off  in 
the  free  state. 

In  the  combination  of  chlorine  (or  oxygen)  and  hydrogen 
chemical  energy  can  also  be  transformed  into  electricity.  To 
accomplish  this,  two  tubes  sealed  at  the  top  and  fitted  there  with 
platinum  electrodes,  reaching  almost  to  the  open  end  of  the  tubes, 
are  filled,  one  with  hydrogen  and  the  other  with  chlorine  (oxygen) 


277.1  ELECTROCHEMISTRY.  425 

and  inverted  in  dilute  sulphuric  acid.  On  connecting  the  electrodes 
by  a  wire  a  strong  current  traverses  the  circuit.  The  gases  ab- 
sorbed in  the  platinum  electrodes  drive  their  ions  into  the  sur- 
rounding liquid,  making  the  H-electrode  negative  and  the  C1(O)- 
electrode  positive.  The  ions  of  hydrogen  and  chlorine  dissolve 
in  the  dilute  sulphuric  acid,  however.  This  apparatus  is  called 
GROVE'S  gas  battery  and  was  known  long  before  a  satisfac- 
tory explanation  of  it  could  be  given. 

It  is  characteristic  of  all  these  various  cells  that  the  reacting 
substances  are  apart  from  each  other.  In  the  oxidation  of  stan- 
nous  chloride  by  mercuric  chloride  the  latter  was  not  put  in  the 
vessel  with  the  stannous  chloride  but  in  the  other  vessel;  in  the 
precipitation  of  silver  chloride  the  silver  nitrate  was  not  put  with  the 
sodium  chloride  solution  but  with  the  sodium  nitrate  solution,  and 
so  on.  The  reaction  took  place  only  because  one  sort  of  ions 
transferred  their  electrification  wholly  or  in  part  through  the  wire 
circuit  to  the  other  electrode,  where  it  either  converted  atoms  into 
ions  or  raised  existing  ions  to  a  higher  potential  or,  possibly,  changed 
ions  of  opposite  potential  sign  to  neutral  atoms. 

Since  we  know  that  chemical  reactions  can  under  suitable  con- 
ditions produce  an  electric  current,  we  can,  conversely,  regard  the 
existence  of  such  a  current  as  an  indication  of  the  occurrence  of 
a  chemical  reaction.  COHEN  has  made  use  of  this  fact  in  deter- 
mining electrically  the  transition  points  of  hydrous  salts  and  other 
systems.  Let  us  take,  for  example,  a  salt  which  loses  its  water 
of  crystallization  at  a  definite  temperature,  e.g.  Glauber's  salt, 
Na2SO4-10H2O;  this  has  a  transition  point  at  about  33°,  where 
the  anhydrous  salt  becomes  capable  of  permanent  existence. 
Now  it  is  possible  for  the  anhydrous  salt  to  remain  in  contact 
with  its  saturated  solution  in  an  unstable  condition  after  the  system 
has  been  cooled  a  few  degrees  below  33°;  the  reverse  is  also  true  of 
the  hydrous  salt.  Since  these  solutions  are  in  contact  with  different 
solid  phases  (one  with  Na2SO4-  10H2O  and  the  other  with  Na2S04) 
they  do  not  have  the  same  concentration;  at  the  transition  point, 
however,  these  concentrations  become  equal,  for  since  both  solid 
phases  are  in  contact  with  the  solution  in  each  case,  the  solubility 
becomes  the  same.  In  his  electrical  method  COHEN  uses  the 
difference  in  concentration  of  the  solutions  which  are  saturated 
in  respect  to  the  two  solid  phases  to  form  a  galvanic  cell.  This 


426 


INORGANIC  CHEMISTRY. 


[§§  277- 


can  be  done  as  follows:  In  the  bottom  of  each  of  the  cylinders 
A  and  B  (Fig.  71)  there  is  a  little  mercury.  A  platinum  wire  is 
fused  into  each  cylinder  and  the  two  are  connected  by  means  of  a 
metallic  wire.  On  top  of  the  mercury  is  some  insoluble  mercurous 
sulphate;  above  this  in  A  is  a  paste  of  Na2SO4-10H2O  and  water, 
in  B  is  a  similar  mixture  of  water  and  Na2SO4.  Below  the  transi- 
tion point  the  solution  in  B  is 
in  the  unstable  condition  and 
more  concentrated  than  that 
in  A,  which  is  stable.  The 
result  is  that  sodium  ions  dif- 
fuse through  the  siphon  from 
the  concentrated  to  the  dilute 
solution,  while  at  the  same 
time  an  equivalent  amount  of 
S04-ions  in  B  combines  with 
part  of  the  mercury  to  form 
mercurous  sulphate,  the  nega- 
tive charge  of  the  sulphate 
ions  being  transferred  to  the 
remaining  mercury.  Thus  an 
electric  current  is  produced 
which  passes  through  the  wire  circuit  from  the  dilute  to  the  con- 
centrated solution.  Its  direction  and  intensity  can  be  determined 
by  inserting  a  galvanometer  in  the  circuit. 

Now,  suppose  that  the  whole  apparatus  is  gradually  warmed; 
the  concentrations  in  A  and  B  will  approach  each  other  as  the 
temperature  nears  the  transition  point  and  at  this  point  they  will 
become  equal.  The  intensity  of  the  current  will  therefore  decrease 
steadily  till  the  transition  point  is  reached,  when  it  is  zero.  If 
the  temperature  is  raised  still  higher  the  solution  in  A  will  become 
unstable  and  more  concentrated  than  that  in  B,  which  latter  will 
then  be  the  stable  solution;  as  a  result  the  direction  of  the  current 
will  be  reversed.  In  this  way  it  is  possible  to  determine  the 
transition  point  very  accurately. 

278.  As  was  remarked  in  §  276,  the  electromotive  fcrce  which 
can  be  obtained  from  chemical  reactions  depends  in  large  jp^asure 
on  the  solution  tensions  of  the  metals.  A  knowledge  of  the  latter 


FIG.  71. 


278.] 


ELECTROCHEMISTRY. 


427 


is  therefore  of  very  great  importance.  It  can  be  acquired  with 
the  aid  of  the  equation  previously  given: 

E  =  IO-4  —  log  -. 
n      *  p 

E,  the  difference  of  potential  between  a  metal  and  the  aqueous 
solution  of  one  of  its  salts  can  be  measured.  All  the  other  quan- 
tities of  this  equation  are  known  with  the  exception  of  P,  which 
can  therefore  be  calculated,  as  is  illustrated  in  the  following  ex- 
ample : 

The  potential  difference  between  magnesium  and  the  normal 
solution  (J  mole  per  liter)  of  its  sulphate  was  determined  to  be 
1.22  volts.  The  equation  thus  becomes: 

1.22=  W~*T  log— , 

since  for  magnesium  n  =  2.  p  is  the  osmotic  pressure  of  the  Mg- 
ions.  On  the  assumption  that  the  salt  is  entirely  split  up  into  ions, 
p  is  22.4  Atm.,  for,  osmotic  pressure  being  equal  to  gas  pressure,  1 
mole  gas  at  0°  and  760  mm.  occupies  a  volume  of  22.4  liters  (§  34) ; 
hence,  if  the  volume  is  a  liter,  the  pressure  becomes  22.4  Atm. 
Therefore,  at  0°  we  have: 

1 .22  =  10~4  X  273(log  P-  log  22.4), 
or  log  P= 43.23,  whence  we  have,  approximately, 

P=  1043. 

The  following  brief  table  indicates  some  of  the  results  for  different 
metals: 


Metal. 

Valence. 

Solution  Ten- 
sion in 
Atmospheres. 

Magnesium        .... 

2 

1043 

Zinc        

2 

1018 

3 

10  13 

Cadmium 

2 

107 

Iron 

2 

103 

\ickel 

2 

10° 

Lead           

2 

10~2 

Hydrogen  

1 
2 

io-4 

10-12 

Mercury 

1 

io-15 

Silver     

1 

io-15 

428 


INORGANIC  CHEMISTRY. 


[§§278- 


The  above  figures  show  how  enormously  the  solution  tension 
differs  in  different  substances.  For  magnesium  and  zinc  it  is 
many  millions  of  atmospheres,  for  copper,  mercury  (ous)  and 
silver  extremely  small  fractions  of  an  atmosphere.  Despite  the 
comparatively  large  errors  in  the  above  data,  due  to  the  difficulty 
of  determining  the  potential  difference  between  the  metal  and  its 
salt  solution,  the  o  r  d  e  r  of  the  decimal  expressing  the  value  of  P 
can  be  accepted  as  reliable  in  each  instance. 

Some  of  these  differences  of  potential  between  metals  and  their 
normal  salt  solutions  are  as  follows: 


Metal. 

Volt. 

Mg 

+  1.22 

Mn        

+  0.798 

Zn  

+  0.439 

Al 

+  0.22 

Cd                ... 

+  0.143 

Fe           

+  0.063 

Tl  

+  0.045 

Co 

-0  043 

Ni             

-  0  .  049 

Pb        

-0.129 

H      

-0.277 

Cu  

-0.606 

Hg 

-  1  .  027 

~is  

\cr 

-  1  .  048 

The  algebraic  sign  of  these  differences  of  potential  can  be 
directly  determined  from  the  solution  tensions.  The  electrolyte 
in  which  zinc  is  immersed  must  assume  a  positive  potential  and 
the  metal  itself  a  negative  potential,  because  no  zinc  solution  can 
be  concentrated  enough  to  hinder  the  emission  of  (positive)  zinc 
ions  by  the  metal.  On  the  other  hand,  copper  must  become  posi- 
tive in  respect  to  a  copper  solution,  for  even  in  the  most  dilute 
solutions  the  osmotic  pressure  of  the  copper  ions  is  greater  than 
the  solution  tension  of  the  metal. 

279.  A  knowledge  of  the  electrochemical  series  of  the  metals 
in  electrolytes  is  of  great  practical  value.  Wherever  combinations 
of  various  metals,  alloys,  metallic  crustations,  etc.,  are  exposed  to 
atmospheric  action  there  is  an  opportunity  to  form  cells  of  short 
circuit.  In  general,  the  metal  with  the  greatest  solution  tension 
goes  into  solution  and  the  other  remains  intact.  A  piece  of  gal- 
vanized (zinc-plated)  iron  wire  does  not  rust,  even  in  those  places 
where  the  plate  has  been  worn  off,  as  much  as  if  it  were  not  zinc- 


280.]  ELECTROCHEMISTRY.  429 

plated.  The  reverse  phenomenon,  that  tinned  iron  rusts  faster 
than  iron  alone,  is  also  due  to  galvanic  causes.  If  our  hypoth- 
esis is  correct  the  atmospheric  moisture  adhering  to  the  metal 
must  act  as  an  electrolyte  with  the  combination  tin-iron  in  such  a 
way  that  iron  becomes  the  dissolving  (negative)  electrode.  Iron 
salts  must  therefore  be  formed  and  then  transformed  into  rust  by 
the  loss  of  carbonic  acid.  The  following  experiment  confirms  this 
view.  Rods  of  iron  and  tin  are  brought  in  contact  by  a  wire  which 
connects  with  a  galvanometer.  If  the  metals  are  dipped  in  water, 
into  which  air  and  carbonic  acid  are  passed  and  to  which  is  added  a 
trace  of  sodium  chloride  (which  always  floats  in  the  air  and  is 
washed  down  by  the  rain),  the  needle  is  deflected.  The  iron  is 
found  to  be  the  anode,  and  in  the  course  of  an  hour  a  thin  yellow 
coating  of  rust  is  to  be  observed  on  it.  Sheet  iron  is  tinned,  as  is 
well  known,  to  prevent  it  from  rusting  (§  199).  If  the  tin  plating 
is  scratched  off  at  any  place  so  as  to  expose  the  iron,  the  latter  begins 
to  rust  very  rapidly,  more  so  even  than  if  it  were  not  tinned. 
Galvanized  iron,  however,  does  not  show  a  trace  of  rust  where  the 
plating  has  been  damaged. 

280.  An  ion  can  only  go  out  of  solution  when  a  force  greater 
than  the  solution  tension  acts  on  it,  just  as  electrically  neutral 
molecules  cannot  crystallize  out  of  a  solution  until  its  osmotic 
pressure  exceeds  that  of  the  saturated  solution.  The  removal  of 
an  ion  can  be  brought  about  by  the  action  of  an  electrical  force. 
This  is  the  real  principle  of  electrolysis.  The  separation  of  an 
ion  from  a  solution  thus  requires  a  definite  electromotive  force, 

2T        P 

which  must  be  equivalent  to  10~4  —  log—  (see  above)  and  must 

therefore  be  stronger  as  the  solution  tension  is  greater  and  the 
osmotic  pressure  of  the  ions  smaller.  But  since  electrolysis  takes 
place  simultaneously  at  both  the  anode  and  the  cathode,  the  total 
force  E  which  is  necessary  for  an  electrolysis  can  be  found  by  taking 
the  sum  of  the  forces  necessary  for  the  separation  of  the  cation  and 
the  separation  of  the  anion,  thus: 


Since  it  is  always  the  case  that  various  sorts  of  anions  and  cations 
are  present  together  in  a  solution,  electrolysis  can  thus  take  place 


430 


INORGANIC   CHEMISTRY. 


[§§280- 


when  E  has  become  large  enough  to  separate  out  one  of  the  varie- 
ties of  cations  and  one  of  the  varieties  of  anions  present. 

This  is  the  basis  of  a  method  of  utilizing  various  electromotive 
forces  to  effect  an  electrolytic  separation  of  metals.  It  is  not  the 
current  strength  which  is  of  primary  importance  to  the  electrolytic 
process  (as  was  formerly  supposed),  but  the  difference  of  potential 
between  the  electrodes.  A  very  successful  example  of  this  method 
is  the  separation  of  copper  from  zinc.  With  a  current  of  low 
voltage  it  is  possible  to  precipitate  only  the  copper  from  a  solution 
containing  ions  of  both  metals;  if  the  electromotive  force  is  in- 
creased, zinc  also  is  separated. 

In  many  cases  the  ions  of  the  water  are  more  easily  separated 
out  than  those  of  the  dissolved  electrolyte.  In  the  electrolysis  of 
potassium  hydroxide,  for  example,  OH-ions  are  liberated  at  the 
anode  (they  are  at  once  decomposed,  however,  into  water  and 
oxygen);  at  the  cathode  it  is  not  potassium  ions  but  hydrogen 
ions  (in  spite  of  their  extremely  small  concentration)  which  are 
discharged,  since  the  solution  tension  of  hydrogen  is  much  less 
than  that  of  potassium. 

281.  The  dissociation  tensions  E  for  various  ions  are  given 
below.  The  figures  are  based  on  equivalent  normal  solutions. 

The  dissociation  tension  of  hydrogen  is  fixed  at  zero  in  the 
table.  Inasmuch  as  there  is  always  an  anode  and  a  cathode. 
it  is  possible  to  subtract  from  all  the  values  of  EI  an  arbitrary  but 
constant  amount  and  add  it  to  the  values  of  E^*  without  affecting 
E(  =  Ei+E2).  The  symbol  O"  represents  a  secondary  ionization 
product  of  the  hydroxyl  ion: 


the  existence  of  which,  according  to  NERNST,  we  are  obliged  to 
assume,  although  only  to  an  extremely  small  degree. 
DISSOCIATION  TENSIONS. 


EI  (Cations). 

E2  (Anions). 

Ag-  -0.771 

I'    -0.520 

Cu"-  0.329 

Br'  -0.993 

H-   0.0 

O"   -1.  23 

Pb"+0.148 

Cl'   -1.417 

Cd"+  0.420 

OH'  -1.67 

Zn-  0.770 

S04"  -1.9 

HSO/-2.6 

281.]  ELECTROCHEMISTRY.  '  431 

These  figures  lead  us  to  important  results.  In  the  first  place 
they  enable  us  to  know  at  once  the  dissociation  tension  of  any 
combination  of  ions.  Zinc  bromide,  for  instance,  will  require 
0.993  +  0.770  =  1.763  volts  for  its  electrolysis;  when  the  concen- 
tration of  the  ions  is  normal,  the  electrolysis  of  hydrochloric  acid 
will  require  1.417  +  0  =  1.417  volts,  and  so  on.  It  is  also  obvious 
that  it  must  be  easy  to  separate  silver  from  copper  electrolytically, 
since  the  difference  of  their  dissociation  tensions  is  almost  0.5  volt. 
It  also  appears  theoretically  possible  to  separate  electrolytically 
iodine  from  bromine  and  bromine  from  chlorine. 

The  order  of  the  metals  in  the  above  electrochemical  series  is 
the  same  as  that  in  which  one  metal  is  precipitated  from  its  solu- 
tion by  the  succeeding  ones.  As  soon  as  a  trace  of  the  dissolved 
metal  is  deposited  on  the  other  one,  the  two  metals  form  with  the 
liquid  an  element,  which  electrolyzes  the  surrounding  solution. 
The  formula 

9T        P 

#  =  10-4—  k)g- 

n      &  p 

tells  us,  however,  that  the  values  of  E  depend  not  only  on  the 
solution  tension  but  also  on  the  osmotic  pressure  of  the  cations. 
Very  decided  changes  in  the  concentration  of  the  salt  solution 
would  make  the  order  of  the  metals  a  different  one.  For  instance, 
it  would  be  possible  to  conceive  a  case  in  which  lead  would  not  be 
precipitated  by  cadmium. 

The  electrochemical  series  of  the  anions  also  brings  out  im- 
portant relations.  Bromine  must  quickly  liberate  iodine  from 
iodide  solutions  and  chlorine  quickly  liberate  bromine  from  bromide 
solutions  because  of  the  marked  difference  in  their  dissociation 
tensions.  We  see,  further,  that  chlorine  must  be  able  to  generate 
oxygen  in  acid  solutions,  but  not  so  with  bromine  or  iodine.  It  is 
also  known,  however,  that  the  generation  of  oxygen  by  chlorine 
proceeds  with  extreme  slowness,  in  sharp  contrast  to  the  rapidity 
with  which  chlorine  deprives  bromine  of  its  negative  charge: 

Cl2+2Br/=Br2-f-2Cr'. 

This  is  not  surprising  in  the  light  of  the  above  considerations,  for 
the  chlorine,  in  order  to  enter  the  ionic  condition,  must  -make  use 
of  the  ion  O",  of  which  there  is  only  an  extremely  small  amount 


432  INORGANIC    CHEMISTRY.  [§§281- 

present.  The  hydroxyl  ion  OH',  which  is  present  in  relatively 
much  larger  amount  and  which  after  the  loss  of  its  negative 
charge  would  also  yield  a  quantity  of  oxygen  equal  to  that  of  the 
chlorine,  holds  its  charge  more  than  0.3  volt  firmer  than  the 
chlorine  ion  in  acid  solution. 

The  application  of  electrolysis  to  commercial  processes  is 
referred  to  in  connection  with  the  substances  concerned  (c/. 
§§  223,  226,  242,  2.45,  248,  and  elsewhere). 

BORON. 

282.  This  element  occurs  in  nature  as  sassolite,  H3B03,6orac^e, 
Mg7Cl2B16O3o,  colemanite,  Ca2B60n-5H2O,  and  borax,  Na2B407* 
10H20.  It  can  be  obtained  in  the  elemental  state  by  the  re- 
duction of  boric  anhydride,  B203,  or  borax  by  means  of  magne- 
sium powder.  It  is  prepared  pure  by  subjecting  a  mixture  of 
boron  chloride  and  hydrogen  to  an  arc  discharge  between  two 
boron  or  water-cooled  copper  electrodes,  or  bringing  the  mixture 
in  contact  with  an  electrically  heated  graphite  tube.  The 
element  melts  between  2000°  and  2500°,  but  has  such  a  high  vapor 
tension  that  it  sublimes  rather  easily  as  low  as  1600°.  Its 
hardness  (exceeded  only  by  diamond),  combined  with  its  amor- 
phous structure,  constitutes  a  valuable  mechanical  characteristic. 
The  electrical  resistance  decreases  with  rising  temperature  at 
remarkable  rapidity. 

It  dissolves  in  molten  aluminium  and,  on  cooling  the  melt,  the 
compound  A1B12  crystallizes  out.  Boron  takes  fire  in  fluorine  and 
chlorine,  uniting  with  them  directly.  When  ignited  in  the  air 
it  burns  to  the  oxide  B203.  At  a  very  high  temperature  it  com- 
bines with  nitrogen  to  form  boron  nitride,  BN.  It  reduces  many 
compounds,  such  as  CuO  and  PbO,  and  decomposes  water  at 
red-heat.  Heating  with  nitric  and  sulphuric  acid  converts  it 
into  boric  acid.  It  is  also  attacked  by  boiling  caustic  alkalies 
(like  aluminium) : 

2B  +  2KOH + 2H20  =  2KB02  +  3H2. 

Boron  hydride. — When  boric  anhydride  is  reduced  with  an 
excess  of  magnesium  powder,  magnesium  boride,  Mg3B2,  is  formed. 


283.]  OXYGEN  COMPOUNDS  OF  BORON.  433 

The  latter  on  being  added  to  hydrochloric  acid  generates  an  ill- 
smelling  gas  consisting  of  hydrogen  and  a  little  boron  hydride. 
This  gas  mixture  burns  with  a  green  flame. 


Halogen  Compounds. 

Boron  chloride,  BC13,  can  be  prepared  by  direct  synthesis, 
but,  better,  by  passing  chlorine  over  boron  carbide.  It  boils 
at  17°.  Its  vapor  density  indicates  the  above  formula.  Water 
breaks  it  up  into  hydrochloric  and  boric  acids;  it  was  with  the 
aid  of  this  reaction  that  the  composition  of  the  compound  was 
determined. 

Boron  fluoride,  BF3,  is  formed,  like  silicon  fluoride  (§  193), 
when  the  oxide  is  warmed  with  a  mixture  of  calcium  fluoride  and 
sulphuric  acid: 

B203 + 3CaF2 + 3H2SO4 = 2BF3 + 3CaS04 + 3H2O. 

It  is  a  gas,  of  which  water  dissolves  700-800  volumes.  A  solution 
of  this  concentration  fumes  in  the  air.  On  dilution  boric  acid 
separates  out  after  some  time;  hydrofluoboric  acid,  HF-BF3,  is 
left  in  the  solution.  This  acid  cannot  be  isolated  in  the  free  state 
but  various  salts  of  it  are  known.  It  thus  displays  a  very  close 
analogy  to  silicon  fluoride. 


Oxygen  Compounds  of  Boron. 

Boron  oxide,  B203,  boric  anhydride,  is  obtained  as  a  vitreous 
mass  by  igniting  boric  acid.  It  is  very  hygroscopic  and  is  recon- 
verted by  the  absorbed  water  into  boric  acid.  With  hydrofluoric 
•acid  it  forms  boron  fluoride.  The  oxide  is  volatile  only  at  elevated 
temperatures. 

283.  Boric  acid,  H3B03,  is  found  in  the  volcanic  districts 
of  Tuscany,  where  jets  of  steam  (the  springs  are  called  "  fumaroles  " 
and  the  jets  proper  "  soffioni  ")  containing  a  little  boric  acid  issue 
from  the  earth.  The  steam  is  conducted  into  water,  in  which  the 
boric  acid  is  retained.  When  this  liquid  reaches  a  certain  concen- 
tration, it  is  allowed  to  settle,  whereupon  it  is  piped  off  into  a  very 


434  INORGANIC  CHEMISTRY.  :  [§283. 

long,  flat  leaden  pan,  which  is  warmed  to  about  50-60°  by  other 
soffioni.  At  this  temperature  the  boric  acid  volatilizes  but  very 
little  with  steam  and  when  the  concentration  has  become  great 
enough  it  crystallizes  out.  It  is  purified  by  converting  it  into 
borax,  which  is  recrystallized  and  then  decomposed  by  hydrochloric 
acid,  setting  free  boric  acid. 

Considerable  boric  a.cid  is  also  made  by  decomposing  native 
borates  with  a  strong  mineral  acid. 

The  volatility  of  boric  acid  with  steam  has  for  a  long  time  been 
regarded  as  an  especially  interesting  phenomenon,  because  the  anhy- 
dride B2O3,  into  which  it  is  readily  converted  at  an  elevated  tempera- 
ture, is  only  volatilized  with  extreme  difficulty.  The  question  there- 
fore arises  as  to  the  particular  compound  in  which  boric  acid  exists 
in  solution  and  the  one  in  which  it  escapes  from  solution. 

The  first  point  can  be  settled  by  a  determination  of  the  boiling- 
point  elevation  or  vapor-tension  lowering  of  boric  acid  solutions.  Meas- 
urements of  this  sort  have  shown  that  H3B03  molecules  exist  in  dilute 
solution.  As  the  solution  becomes  more  concentrated  the  vapor-tension 
lowering  no  longer  corresponds  to  this  formula;  the  decrease  in  the 
lowering  indicates  that  the  number  of  molecules  of  dissolved  substance 
has  grown  less,  i.e.  some  such  change  as  4H3BO3=H2B4O7  +  5H20  has 
occurred. 

If  H3BO3  molecules  volatilize  with  the  water,  the  concentrations  of 
boric  acid  in  the  solution  and  in  the  vapor  must,  according  to  HENRY'S 
law,  maintain  a  constant  ratio,  independent  of  the  amount  of  boric 
acid  present.  This  was  found  to  be  true  for  dilute  solutions  but  not 
for  concentrated  ones,  which  is  in  agreement  with  the  experiments  on 
vapor-tension  lowering,  because  there  also  the  concentration  of  the  acid 
in  the  vapor  remained  proportional  to  its  concentration  In  the  solution. 
It  is  therefore  demonstrated  that  the  compound  which  escapes  with 
the  steam  is  boric  acid,  H3B03. 

Boric  acid  crystallizes  in  lustrous  laminse,  which  feel  greasy  and 
are  difficultly  soluble  in  cold  water  (about  3%  at  ordinary  tem- 
peratures). -This  solution  acts  as  a  weak  antiseptic,  for  which 
purpose  it  is  frequently  used.  At  100°  boric  acid  loses  1  molecule 
H20,  passing  over  into  metaboric  acid,  HB02.  At  140°  tetraboric 
acid,  H2B407  (=4B(OH)3~5H20),  is  formed,  the  sodium  salt  of 
which  is  borax. 

No  salts  of  the  normal  boric  acid,  B(OH)3,  are  known,  but 


§  283.]  BORON.  435 

metaboric  acid  forms  several.     They  are  unstable  and  are  converted 
by  carbon  dioxide  into  salts  of  tetraboric  acid: 

4NaB02  +  CO2  =  Na2B407+  Na2C03. 

The  best-known  salt  of  boric  acid  is  borax,  Na2B407-12H2O, 
often  called  tinkal.  At  present  most  of  the  borax  on  the  market 
is  made  by  boiling  colemanite,  Ca2B6On-5H2O,  found  in  California, 
or  a  similar  borate,  occurring  in  Chile,  with  soda.  Borax  swells 
greatly  on  heating;  this  is  due  to  the  escape  of  water  of  crystalliza- 
tion from  the  semi-molten  salt.  On  continued  heating  it  forms  a 
vitreous  mass.  This  glass  has  the  property  of  dissolving  metallic 
oxides,  some  of  which  give  double  borates  of  a  characteristic  color; 
hence  its  use  in  qualitative  analysis.  The  same  property  makes 
it  valuable  in  soldering;  solder  adheres  only  to  the  untar- 
nished metal,  so  a  little  borax  is  placed  on  the  surface  of  the  metal 
and  heated  with  the  soldering-iron  in  order  to  remove  the  rust. 
The  dissolving  of  metallic  oxides  is  easily  understood,  when  we 
write  Na2B4O7  as  2NaB02+B203;  it  is  the  boric  oxide,  B203, 
which  can  be  regarded  as  combining  with  the  metallic  oxides  to 
form  salts. 

Boric  acid  is  a  weak  acid;  its  salts  are  therefore  hydrolyzed 
quite  perceptibly — more  so,  of  course,  as  the  dilution  increases. 
This  can  be  illustrated  by  a  simple  experiment  devised  many  years 
ago  by  ROSE.  To  a  concentrated  solution  of  borax  some  litmus 
is  added  and  then  acetic  acid  until  the  litmus  is  just  red;  if  the 
liquid  is  then  diluted,  it  turns  blue  because  the  alkali  is  set  free 
and  boric  acid  has  scarcely  any  effect  on  litmus. 

Rather  interesting,  also,  is  the  behavior  of  silver  borate,  which 
is  deposited  as  a  white  salt  on  mixing  concentrated  solutions  of 
borax  and  silver  nitrate.  When  dilute  solutions  are  mixed,  how- 
ever, a  precipitate  of  grayish-brown  silver  oxide  is  formed,  the 
silver  borate  being  almost  completely  hydrolyzed  in  the  dilute 
solution. 

On  treating  a  mixture  of  boric  acid  and  sodium  peroxida 
with  water  a  perborate  is  formed  and  crystals  of  the  composi- 
tion NaB03-4H2O  separate  out.  They  are  stable  when  solid 
but  liberate  oxygen  from  a  warm  solution.  The  solution  con- 
tains hydrogen  peroxide  also. 


436  INORGANIC  CHEMISTRY.  [§§284- 


ALUMINIUM. 

284.  This  metal  does  not  occur  native,  but  in  combination  it 
is  found  in  large  quantities  and  very  widely  diffused.  Corundum, 
including  the  precious  stones  sapphire  and  oriental  ruby  and  the 
natural  abrasive  emery  (all  noted  for  their  hardness),  consists  of 
alumina  A12O3,  colored  by  traces  of  other  oxides.  Bauxite  is  a 
hydrate  of  aluminium  and  iron.  Clay  and  kaolin  (China  clay)  are 
principally  aluminium  silicate.  Many  other  minerals,  such  as 
feldspar,  mica,  etc.,  contain  it  as  a  base.  A  peculiar  aluminium 
mineral,  cryolite  or  ice  stone,  3NaF-AlF3,  is  found  in  Greenland. 

The  metal  can  be  obtained  from  the  chloride  by  reduction 
with  sodium  but  at  present  it  is  produced  exclusively  by  decom- 
posing aluminium  oxide  with  the  electric  current. 

The  most  important  commercial  process  is  that  of  HALL 
(invented  independently  in  Europe  by  HEROULT).  Alumina  is 
dissolved  in  a  fused  bath  consisting  of  cryolite  or  an  equivalent 
mixture.  The  process  is  carried  out  in  a  large  carbon-lined  pot, 
the  inner  surface  of  which  constitutes  the  cathode.  Carbon  rods 
immersed  in  the  bath  serve  as  anodes.  Fresh  alumina  is  added 
from  time  to  time  and  the  metal  is  drawn  off  at  the  bottom 
periodically.  The  temperature  is  a  little  above  the  melting-point 
of  cryolite.  A  current  of  several  thousand  amperes  and  less  than 
8  volts  maintains  the  liquidity  of  the  bath  as  well  as  effects  the 
electrolysis. 

The  increased  output  due  to  improved  methods  has  brought 
the  price  of  the  metal  down  from  over  $90  per  pound  in  1856  to 
about  $0.20  at  the  present  time,  and  the  production  is  steadily 
increasing  (34,000,000  Ibs.  in  1909). 

Aluminium  is  a  silvery-white  metal  of  low  specific  gravity 
(2.583).  It  is  rigid  but  very  ductile  and  malleable.  It  softens 
at  about  600°,  melts  at  658°,  and  boils  at  about  1800°. 

It  is  permanent  in  the  air,  since  it  soon  becomes  coated  with  a 
firm  thin  layer  of  oxide.  Small  fragments  burn  with  a  bright 
light  when  heated  in  an  oxygen  atmosphere.  It  is  not  attacked 
by  dilute  nitric  acid  at  ordinary  temperatures  and  only  slightly  so 
by  dilute  sulphuric  acid.  Hydrochloric  acid  dissolves  it  readily, 


285.]  ALUMINIUM.  437 

as  does  also  caustic  potash,  hydrogen  being  evolved  and  a  1  u  m  i  - 
nates  formed  in  the  latter  case. 

Various  alloys  of  aluminium  have  found  a  place  in  the  arts.  Among 
them  mention  may  be  made  of  aluminium  bronze,  which 
consists  of  copper  and  5-12%  aluminium.  It  can  be  easily  cast  and 
has  a  golden  color  and  lustre.  Its  great  firmness  and  elasticity  render 
it  valuable  for  physical  instruments  (balance  beams)  and  watch  springs. 
New  alloys  of  aluminium  are  being  constantly  brought  on  the  market: 
there  is  one  with  magnesium  called  magnalium  and  another  with  tung- 
sten, for  example. 

Aluminium  reduces  many  oxides  (GOLDSCHMIDT)  with  a  vigor- 
ous evolution  of  heat  (§  293).  The  reduction  proceeds  of  itself 
after  it  has  been  started  at  a  certain  place  in  the  mixture.  For  this 
purpose  a  primer  is  used  consisting  of  a  mixture  of  oxygen-producing 
substances,  such  as  KClOs,  etc.,  and  a  piece  of  magnesium  ribbon, 
which  is  ignited  with  a  match.  The  heat  that  is  thus  evolved  is 
used  to  heat  iron  bolts  to  white-heat  and  also  for  welding  railroacj 
rails,  etc.  The  welding  is  accomplished  by  packing  the  rails  in  a 
mixture  of  iron  oxide,  sand,  and  aluminium  powder  together 
with  a  special  sort  of  cement  for  making  it  compact.  When 
this  mass  is  ignited  it  continues  to  burn  and  heats  the  rails  to 
glowing. 

An  amalgam  of  aluminium  is  easily  prepared  by  introducing 
aluminium  filings  into  a  J%  solution  of  corrosive  sublimate.  This 
amalgam  decomposes  water  energetically  at  ordinary  tempera- 
tures, liberating  hydrogen  and  forming  aluminium  hydroxide.  As 
neither  basic  nor  ad^d  substances  go  into  solution,  it  is  a  neutral 
reducing-agent.  The  cause  of  this  energetic  reaction  is  due  to  the 
circumstance  that  the  mercury  hinders  the  formation  of  a  thin 
firm  coating  of  oxide  over  the  surface  of  the  metal,  which  would 
otherwise  protect  it  from  further  oxidation. 

Compounds  of  Aluminium. 

285.  The  only  known  oxide  of  aluminium  is  alumina,  A1203, 
which  is  formed  on  heating  aluminium  salts  or  the  hydroxide.  It 
is  a  white  amorphous  powder,  readily  soluble  in  acids;  however, 
after  it  has  been  strongly  ignited  it  is  no  longer  soluble  and  must 


438  INORGANIC  CHEMISTRY.  [§  285. 

then  be   disintegrated    by  fusion  with  potassium  hydroxide  or 
acid    potassium  sulphate.     It  is   found   crystallized    in  nature 

(§284). 

The  artificial  manufacture  of  rubies  and  sapphires  is  accomplished 
by  fusing  amorphous  A12O3  with  lead  oxide  at  bright  red  heat  in  a 
Hessian  crucible.  Lead  aluminate  is  first  formed,  whereupon  the  silica 
of  the  crucible  causes  the  alumina  to  separate  out  in  beautiful  crystals, 
exactly  like  the  natural  gems.  By  adding  a  little  potassium  dichro- 
mate  we  get  crystals  having  the  color  of  the  natural  rubies;  similarly, 
the  addition  of  cobalt  oxide  gives  sapphires. 

Aluminium  hydroxide,  Al203-nH20,  is  deposited  as  a  hy- 
drogel  (§  195)  when  a  solution  of  an  aluminium  salt  is  treated 
with  ammonia.  In  the  decomposition  of  the  aluminates  it  is 
obtained  as  a  white  powder.  A  hydrate  with  a  low  percentage  of 
water,  A1203-2H20,  bauxite,  occurs  in  France  and  different  parts 
of  the  United  States  in  large  deposits.  Aluminium  hydroxide  is 
both  weakly  acidic  and  weakly  basic  in  character.  Its  salts  with 
acids  suffer  partial  hydrolysis  in  aqueous  solution  and  hence  react 
acid  (§  239).  It  dissolves  in  alkalies  to  form  aluminates,  such  as 
AlC^K,  AlC^Na,  and  AlOsNas,  which  are  deposited  in  the  amor- 
phous state  when  alcohol  is  added  to  their  aqueous  solutions. 
They  are  decomposed  by  atmospheric  carbonic  acid. 

Aluminium  hydroxide  is  insoluble  in  water  but  dissolves  in 
a  solution  of  aluminium  chloride.  By  subjecting  this  solution 
to  dialysis,  it  is  possible  to  get  rid  of  the  hydrochloric  acid 
(which  is  present  because  of-  hydrolytic  dissociation)  entirely 
and  thus  obtain  a  colloidal  solution  of  the  hydroxide.  Alumin- 
ium hydroxide  does  not  form  salts  with  weak  acids. 

Aluminium  chloride,  Aids,  is  most  conveniently  prepared  by 
passing  dry  hydrochloric  acid  gas  over  aluminium  filings  in  a  tube 
of  porcelain  or  glass  and  collecting  the  sublimed  product  in  a 
wide-mouthed  bottle  (see  Fig.  72) .  After  the  tube  has  been  heated 
to  a  sufficiently  high  temperature  to  start  the  reaction,  no  further 
heating  is  required;  however,  it  is  more  practicable  to  continue 
heating  in  order  to  collect  the  chloride  in  the  receiver. 

Aluminium  chloride  is  very  hygroscopic.  The  aqueous  solu- 
tion hydrolyzes  so  readily,  depositing  alumina,  that  it  can  only  be 
preserved  by  the  addition  of  an  excess  of  hydrochloric  acid.  Such 


§  285.]  COMPOUNDS  OF  ALUMINIUM.  439 

a  solution  does  not  yield  aluminium  chloride  on  evaporation,  since 
it  decomposes  completely  into  the  hydroxide  and  hydrochloric 
acid  on  account  of  the  continued  removal  of  the  latter  dissociation 
product.  The  vapor  density  of  the  chloride  up  to  400°  corre- 
sponds to  the  formula  A12C16,  above  760°  to  A1C13.  With  the 
chlorides  of  potassium  and  sodium,  aluminium  chloride  forms 
compounds  such  as  A1C13-KC1,  whose  solutions  can  be  evap- 
orated without  decomposition.  Compounds  such  as  Aids  •  PC13, 


FIG.  72. — PREPARATION  OF  ALUMINIUM  CHLORIDE. 

A1C13-POC13,  etc.,  have  also  been  prepared.  In  organic  chemistry 
anhydrous  aluminium  chloride  is  of  great  value  in  synthetical  work. 
Aluminium  sulphate,  Al2(SO4)3-16H20,is  obtained  by  treating 
clay  with  concentrated  sulphuric  acid;  the  product  is  dissolved  in 
water  and  allowed  to  crystallize.  Aluminium  sulphate  unites  with 
the  alkali  salts  to  form  double  salts  of  the  general  type: 

i  in 

R2SO4  -  R2/(S04)3  •  24H2O, 

which  are  known  as  alums.  R  may  be  either  K,  Na,  NH4,  Cs, 
Rb,  Tl,  or  an  organic  base ;  R'  may  be  Fe  (ic)  or  Cr,  instead  of  Al. 
The  alums  all  crystallize  in  octahedrons  and  cubes,  which  often 
grow  to  large  dimensions;  they  form  mixed  crystals  readily. 
Ordinary  alum  (potassium  alum)  is  used  as  a  mordant  in  dyeing 
fORG.  CHEM.,  §  362),  but  it  is  being  gradually  superseded  as  such 
by  aluminium  sulphate  and  sodium  aluminate.  In  the  vicinity  of 
Rome  the  mineral  alunite,  or  alum  stone,  is  found,  whose  com- 
position is  K(A102H2)3(SO4)2;  from  it  a  much  sought  variety  of 
alum  is  made.  Alum  is  also  made  from  cryolite,  etc. 

When  two  salts  combine  we  may  have  one  of  two  results! 
either  the  new  salt  which  is  formed  gives  ions  hi  dilute  aqueous 


440  INORGANIC  CHEMISTRY.  [§§285- 

solution  that  differ  from  those  of  the  two  salts,  or  it  gives  the  same 
ions.  A  good  example  of  the  former  case  is  yellow  prussiate  of 
potash;  it  gives  neither  ferrous  ions  nor  cyanide  ions,  so  that  it 
must  be  regarded  as  K4[Fe(CN)6].  Such  salts  are  termed  com- 
plex. The  second  case  is  illustrated  by  the  alums.  A  dilute 
alum  solution  exhibits  all  the  reactions  which  characterize  its  com- 
ponents and  its  conductivity  is  the  mean  of  the  two  separate  salts  for 
the  same  concentration.  When  the  union  is  of  this  sort  we  have 
what  is  called  a  double  salt.  Between  the  two  kinds  there  are 
salts  of  an  intermediate  nature  which  form  not  only  complex  ions 
but  also  the  original  ions  to  a  greater  or  less  extent.  The  copper- 
ammonia  compounds  (§  244)  behave  in  this  way. 

286.  Aluminium  silicate,  kaolin,  is  formed  in  nature  by  the 
weathering  of  the  numerous  alkali-alumina  double  silicates,  the 
alkali  silicate  being  dissolved  out,  leaving  the  insoluble  aluminium 
silicate.  Clay  is  aluminium  silicate;  it  is  usually  colored  brown 
by  iron  oxide.  It  is  the  essential  raw  material  of  the  ceramic 
industries,  being  used  both  for  rough  bricks  and  the  finest  china- 
ware;  of  course  the  better  grades  require  better  sorts  of  clay.  Bricks 
are  molded  out  of  ferruginous  and  calcareous  clays  (loam)  and  then 
baked  ("burned,"  or  "fired")  till  they  become  firm.  Under  the 
head  of  earthenware,  or  porous  ware  (faience,  majolica,  etc.,  and 
common  crockery)  we  include  all  articles  which  consist  of  burned 
clay  (frequently  mixed  with  quartz),  are  porous  and  display 
an  earthy  fracture  and  which  are  covered  with  a  glaze  of  easily 
fusible  silicates.  The  glaze  is  produced  by  introducing  salt  into 
the  kiln.  The  hot  steam  causes  the  formation  of  hydrochloric  acid 
and  souum  hydroxide,  which  unite  with  the  clay  to  form  sodium 
aluminium  silicate.  In  porcelain  the  pores  of  the  earthen  mass 
are  completely  filled  with  fused  silicate,  as  a  result  of  the  addition 
of  feldspar  and  quartz  before  the  burning.  The  less  of  such  admix- 
tures is  present  the  more  difficult  the  porcelain  is  to  burn  and  the 
less  sensitive  it  is  to  changes  of  temperature. 

Clay  is  the  most  widely  diffused  refractory  material;  it 
resists  not  only  high  temperatures  and  sudden  changes  of  tem- 
perature., but  chemical  action  as  well. 

Ultramarine  is  a  very  beautiful  blue  pigment,  which  is  prepared  arti- 
ficially by  heating  a  mixture  of  clay,  soda,  sulphur  and  wood  charcoal 
in  the  absence  of  air.  It  occurs  in  nature  as  lapis  lazuli.  It  is  usually 


288.]  GALLIUM,  INDIUM,   THALLIUM.  441 

regarded  as  a  compound  of  sodium  aluminium  silicate  with  polysul- 
phides  of  sodium.  This  is  indicated  by  the  fact  that  it  is  attacked  by 
acids  with  the  evolution  of  hydrogen  sulphide  and  the  disappearance 
of  the  color,  while  it  is  unaffected  by  alkalies.  It  is  still  uncertain  what 
substance  gives  the  pigment  its  blue  color. 


GALLIUM,  INDIUM,  THALLIUM. 

287.  The  existence  of  gallium  was  predicted  by  MENDELEEFF  (§217) 
in  the  same  manner  as  that  of  germanium.     The  hypothetical  eka-alu- 
minium  was  discovered  in  1875  by  LECOQ  DE  BOISBAUDRAN  in  a  zinc 
blende  by  means  of  spectrum  analysis.     Its  spectrum  consists  of  two 
violet  lines.     It  is  a  very  rare  element.     The  metal  is  white,  melts 
as  low  as  30°  and  has  a  specific  gravity  of  5.9.     It  is  only  superficially 
oxidized  by  the  air  and  is  not  attacked  by  water.    Like  aluminium, 
it  is  only  slightly  affected  by  nitric  acid  but  dissolves  readily  in  hydro- 
chloric acid  as  well  as  ammonia  and  potassium  hydroxide.     It  forms 
alloys  with  aluminium,  which,  when  the  proportion  of  aluminium  \-: 
small,  are  liquid  at  ordinary  temperatures   because  of  the  depression 
of  the  melting-point  of  gallium,  and  it  decomposes  water  almost  as 
readily  as  sodium. 

In  its  compounds,  also,  gallium  displays  much  analogy  with  aluminium. 
The  hydroxide  dissolves  in  alkalies.  The  chloride,  GaCl3,  fumes  in  the 
air  like  A1C13  and  its  aqueous  solution  yields  hydrochloric  acid  on  evap- 
oration. The  sulphate  gives  an  alum,  Ga^SC^V  (NH4)2S04-24H2O, 
with  ammonium  sulphate.  Hydrogen  sulphide  precipitates  gallium 
only  from  an  acetic  acid  solution,  in  which  respect  gallium  resembles 
zinc  (§  269). 

Indium  has  already  been  referred  to  in  the  discussion  of  the  periodic 
fcystem  (§216),  so  that  it  will  be  passed  over  here  with  a  brief  descrip- 
tion. It  was  discovered  through  its  spectrum,  a  blue  line.  This  ele- 
ment, too,  occurs  very  rarely,  being  found,  in  certain  blendes.  The 
m  e  t  a  1  is  white;  m.-pt.,  176°;  sp.  g.,  7.42.  It  is  permanent  in  the  air; 
heated  to  a  high  temperature  it  burns  with  a  blue  flame  to  the  oxide 
In2O3.  The  chloride  InCl3  is  hygroscopic;  its  aqueous  solution  does 
not  decompose  on  evaporation.  The  sulphate  forms  an  alum  with 
ammonium  sulphate.  The  hydroxide  dissolves  in  alkalies. 

288.  Thallium  is  the  most  common  of  these  three  elements,  notwith- 
standing it  always  occurs  in  limited  amounts.     It  is  occasionally  found  in 
the  "  Abraum  salts"  carnallite  and  sylvite  and  frequently  also  in  different 
native  sulphides.     When  the  zinc  blendes  are  roasted  in  sulphuric  acid 
factories  the  thallium  goes  off  with  the  fumes  and  settles  in  the  flue  dust 
and  chamber  mud.     From  these  deposits  it  is  obtained  by  boiling  with 


442 


INORGANIC   CHEMISTRY. 


[§§288- 


dilute  sulphuric  acid  and  precipitating  with  hydrochloric  (or  better  hydri- 
odic)  acid,  whereupon  the  sparingly  soluble  chloride  (or  iodide)  is 
deposited.  This  element  was  also  discovered  with  the  spectroscope 
(CROOKES);  its  spectrum  is  a  bright  green  line. 

Thallium  is  a  soft  metal,  about  like  sodium,  and  has  a  bluish  color  like 
lead.  Sp.  g.,  11.8;  m.-pt.,  290°.  In  moist  air  it  oxidizes  very  rapidly 
at  the  surface;  but  it  does  not  decompose  water  at  ordinary  tempera- 
tures. When  heated  it  burns  with  a  beautiful  green  flame.  Sulphuric 
and  nitric  acids  dissolve  it  readily,  but  hydrochloric  acid  acts  very  slowly 
because  of  the  slight  solubility  of  the  chloride. 

There  are  two  sets  of  compounds:  the  thallous  compounds,  derived 
from  the  oxide  T120,  and  the  thallic  compounds,  from  the  oxide  T1203. 
The  former  resemble  those  of  the  alkalies  and  silver  very  much.  This 
similarity  shows  itself,  for  instance,  in  the  solubility  of  the  hydroxide 
and  the  carbonate,  whose  solutions  react  alkaline.  Moreover,  many 
thallium  salts  are  isomorphous  with  potassium  salts  and,  like  the  latter, 
give  double  salts  with  platinum  chloride,  e.g.  Tl2PtCl6.  Further 
there  is  an  alum  T12S04  •Al^SO^^H.jO,  as  well  as  other  double  sul- 
phates, e.g.  Tl2S04-MgSO4-6H20,  which  are  analogous  to  the  corre- 
sponding potassium  double  salts.  On  the  other  hand  thallium  resembles 
silver  and  lead  in  the  small  solubility  of  its  halides  (the  iodide  is  the 
least,  and  the  chloride  the  most,  soluble)  and  also  in  respect  to  the 
order  of  solubility  of  these  compounds. 

In  the  thallic  compounds  the  element  is  trivalent,  like  the  other 
elements  of  the  group;  furthermore,  like  the  compounds  of  the  latter, 
the  thallic  compounds  readily  form  complex  salts,  and  undergo  con- 
siderable hydrolysis  when  dissolved  in  water. 


SUMMARY  OF  THE  GROUP. 

289.  The  five  elements  last  considered,  B,  Al,  Ga,  In,  Tl,  form 
a  natural  group,  in  which  the  last  three  display  particular  similarity 
to  each  other  in  their  physical  properties.  Something  analogous 
was  observed  with  copper,  silver  and  gold  in  the  first  group  and 
with  zinc,  cadmium  and  mercury  in  the  second  group.  The  fol- 
lowing table  affords  a  brief  comparison  of  certain  physical  data: 


B 

Al 

Ga 

In 

Tl 

11  0 

27  1 

69  9 

114  8 

204  0 

Specific  gravity  
Melting-point  

2.45 
2000° 

2.58 
658° 

5.9 
30° 

7.4 
176° 

11.8 
290° 

290.]  THE   RARE  EARTHS.  443 

In  the  spectra  of  Ga,  In  and  Tl  it  is  again  noticeable  that  the 
lines  move  towards  the  red  end  as  the  atomic  weight  increases 
(§  265). 

As  to  their  chemical  nature  it  may  be  remarked  that  all  the 
elements  of  this  group  are  trivalent  and  that  the  basicity  of  their 
oxides  increases  with  rising  atomic  weight;  boron  hydroxide  (boric 
acid)  has  exclusively  acid  properties,  but  the  hydroxides  of  the 
other  elements,  even  TT^OH)3;  are  also  soluble  in  alkalies.  As 
most  of  the  lower  oxides  of  the  metals  are  more  strongly  basic  than 
the  higher  oxides,  it  is  not  strange  that  thallous  hydroxide  is  a  strong 
base. 

THE  RARE  EARTHS. 

290.  In  the  middle  of  the  periodic  table  (p.  301)  are  located  a  number 
of  elements,  which  are  classed  under  the  term  "rare  earths/'  There  is 
still  much  uncertainty  in  regard  to  some  of  them,  particularly  as  to 
their  elemental  nature.  This  is  due  in  large  measure  to  the  great  simi- 
larity between  ihe  elements  and  the  consequent  difficulty  in  separating 
them.  They  may  be  arranged  in  two  groups :  the  cerium  group  con- 
taining the  elements  lanthanum,  cerium,  praseodymium,  neodymium 
and  samarium  ;  and  the  yttrium  group  containing  europium,  terbium, 
dysprosium,  holmium,  yttrium,  gadolinium,  erbium,  thulium,  ytter- 
bium, scandium,  and  lutecium. 

These  elements  occur  in  various  rare  minerals  which  have  been  found 
principally  in  Sweden  and  Greenland,  viz.,  cerite,  gadolinite,  euxenite, 
orthite,  etc. 

Since  the  use  of  the  oxides  of  cerium  and  thorium  in  the  incandescent 
gas-light  of  AUER  VON  WELSBACH  has  created  a  demand  for  them,  minerals 
in  which  the  rare  earths  occur  are  being  ardently  sought.  The  interesting 
fact  has  developed  that  they  are  by  no  means  so  "rare  "  as  was  supposed. 
An  especially  rich  source  of  these  earths  has  been  found  in  monazite  sand, 
which  occurs  in  rather  large  quantities  in  the  United  States  (production 
900,000  Ibs.  annually) ,  Canada  and  Brazil.  It  consists  chiefly  of  a  phos- 
phate of  Ca,  La,  Di,  Y  and  Er,  with  varying  amounts  of  thorium  silicate 
and  thorium  phosphate. 

In  order  to  isolate  the  rare  earths  from  these  minerals  the  latter  are 
powdered  very  finely  and  heated  to  faint-red  heat  with  concentrated  sul- 
phuric acid.  Thus  the  rare  earths  are  changed  into  sulphates  and  the 
silicic  acid  is  converted  into  the  insoluble  condition.  The  sulphates  are 
then  taken  up  in  ice-water,  in  which  they  dissolve  much  more  readily 
than  in  warm  water  (since  a  difficultly  soluble  hydrate  is  formed  at  a 


444  INORGANIC    CHEMISTRY.  [§290. 

higher  temperature).  From  this  cold  solution  they  can  be  precipitated 
with  oxalic  acid,  their  oxalates  being  almost  insoluble  even  in  dilute 
acids.  Thus  they  are  freed  from  Ca,  Fe,  etc.  The  oxalates  are  then 
converted  into  oxides  by  heating. 

The  separation  of  these  oxides  is  a  more  difficult  task.  Various 
methods  are  in  use,  by  which  the  separation  of  the  eerie  earths  is 
fairly  well  accomplished  ;  but  for  the  numerous  ytteric  earths  no  suc- 
cessful method  has  yet  been  devised.  Some  of  the  methods  employed 
are  as  follows:  The  insolubility  of  the  sulphates  of  cerium,  lanthanum 
and  didymium  in  a  saturated  sodium  sulphate  solution  (by  reason  of 
the  formation  of  double  salts)  is  made  use.  of  to  separate  them  from 
erbium,  ytterbium  and  yttrium.  The  nitrates  of  the  various  metals  of 
this  group  differ  markedly  in  their  stability  on  heating;  hence  another 
method  of  separation  has  been  devised,  by  which  the  nitrates  are  decom- 
posed one  after  another  by  heating  and  those  that  remain  undecom- 
posed  after  each  successive  heating  are  extracted  with  wrater.  A  third 
method  is  the  fractional  precipitation  of  the  solutions  with  ammonia. 
Further,  by  fractional  precipitation  with  potassium  chromate  (the 
insoluble  neutral  chromates  being  deposited),  separations  can  be  accom- 
plished which  are  otherwise  very  difficult. 

The  ytteric  earths  are  separated  after  URBAIN  by  mixing  the 
aqueous  solution  of  their  nitrates  with  bismuth  nitrate,  which  is  iso- 
morphous  with  them.  From  this  solution  they  separate  out  as  mixed 
crystals  with  the  bismuth  nitrate,  the  different  kinds  of  mixed  crystals 
having  different  solubilities.  If  a  fractional  crystallization  is  con- 
ducted, certain  of  the  rare  earths  accumulate  in  the  first  crystalliza- 
tions, others  in  the  last  crystallizations,  the  middle  ones  consisting 
almost  wholly  of  bismuth  nitrate. 

Most  of  the  rare  earth  metals  form  only  one  oxide,  having  the 
formula  M2O3;  cerium,  praseodymium  and  neodymium  have  higher 
oxides  as  well ;  of  these  Ce02  is  able  to  form  salts. 

The  best  method  of  detecting  these  metals  is  by  spectroscopy. 
The  spectra  of  the  eerie  metals  are  satisfactorily  known ;  those  of  the 
ytteric  metals  are  not  so  well  known.  Many  of  the  latter  are  characterized 
by  absorption  bands,  thus  dysprosium,  holmium  and  thulium.  Others, 
like  yttrium  gadolinium  and  ytterbium,  whose  oxides  and  salts  are 
colorless,  do  not  give  an  absorption  spectrum,  but  their  spark  spectrum 
is  characteristic.  The  spectra  of  the  ytteric  earths  display  a  great 
many  lines.  Furthermore,  investigations  of  the  ultraviolet  spectra 
(photographic)  have  furnished  important  information. 

In  addition  to  these  kinds  of  spectra  the  phosphorescence  spectrum 
should  be  mentioned  as  an  important  means  of  investigation,  especially 


THE  RARE  EARTHS.  445 

to  determine  the  purity  of  these  earths.  When  the  earths  are  placed 
in  an  evacuated  tube  and  exposed  to  the  action  of  cathode  rays,  the 
earths  become  luminous,  a  phenomenon  that  is  known  as  cathodic 
phosphorescence.  The  spectrum  of  the  phosphorescence  has  character- 
istic lines.  It  has  been  proved  that  the  perfectly  pure  earths  do  not 
show  the  phospherescence,  but  that  it  is  caused  by  extremely  slight 
admixtures  of  other  earths.  The  maximum  influence  is  caused  in  most 
cases  by  an  admixture  of  1-0.1%  The  disappearance  of  the  phosphor- 
escence is  therefore  a  means  of  telling  when  the  earth  is  pure.  On 
the  other  hand,  however,  the  characteristic  phosphorescence  spectrum 
can  be  used  to  recognize  some  of  the  earths. 

Cerium  occurs  principally  in  cerite  (as  high  as  60%).  Its  salts  are 
colorless  when  pure  and  give  no  absorption  spectrum  (§  263). 

The  metal  looks  like  iron  but  is  soft,  like  lead.  It  oxidizes  slowly 
in  the  air,  becoming  coated  with  a  black  layer.  At  an  elevated  tem- 
perature it  takes  fire.  An  alloy  of  70%  Ce  and  30%  Fe  gives  off  sparks 
when  it  is  scratched,  and  it  is  sometimes  used  as  a  substitute  for  matches. 
It  forms  two  sets  of  salts,  the  cerous  salts,  which  can  be  derived  from 
the  oxide  Ce203  and  are  colorless,  and  the  eerie  salts,  derivable  from 
Ce02,  which  are  yellow  or  brown.  Cerium  is  thus  quadrivalent  (as  the 
existence  of  the  fluoride  CeF4-H2O  also  indicates)  and  so  belongs  to  the 
fourth  group  of  the  periodic  system.  When  chlorine  is  passed  into  an 
alkaline  solution  of  a  cerous  salt  a  yellow  precipitate  of  CeO2  is  obtained. 

Lanthanum  is  only  trivalent.  Its  oxide  La2O3  and  its  salts  are  color- 
less when  pure. 

Didymium  was  formerly  regarded  as  an  element,  but  AUER  VON  WELS- 
BACH  succeeded  in  splitting  it  up  into  two  components,  called  praseody- 
mium and  neodymium.  This  can  be  accomplished  by  making  use  of 
the  difference  in  solubility  of  their  potassium  double  sulphates  in  a 
concentrated  solution  of  potassium  sulphate.  The  praseodymium  salts 
are  green  and  give  green  solutions;  the  neodymium  salts  have  an  ame- 
thyst color  and  give  pink  solutions.  The  absorption  spectra  of  the 
two  elements  differ  considerably. 

Scandium  occurs  in  the  mineral  wolframite.  It  is  a  trivalent  element, 
like  La.  Its  existence  was  predicted  by  MENDELEEFF*  who  called  it 
ekaboron.  Its  tri valence  places  it  in  the  aluminium  group.  The  hydrox- 
ide Sc(OH)3  is  gelatinous,  but  insoluble  in  an  excess  of  alkali.  Pure 
scandium  can  be  prepared  by  way  of  the  sodium  double  carbonate, 
Sc2(CO3)3-4Na2C03-6H20. 

Ytterbium.  —  The  oxide  Yb2O3  is  the  main  constituent  of  erbia, 
formerly  regarded  as  an  elementary  oxide  (obtained  from  euxenite  and 
gadolinite),  but  nowr  known  to  contain  also  the  oxides  of  scandium, 


446  INORGANIC  CHEMISTRY.  [§§290- 

yttrium,  erbium,  etc.  Ytterbia  (oxide)  is  obtained  by  fractional  heating 
of  the  nitrate  mixture  (see  above).  The  salts  of  ytterbium  are  colorless 
and  give  no  absorption  spectrum.  AUER  VON  WELSBACH  recently  suc- 
ceeded in  resolving  ytterbium  into  two  elements,  which  he  called  alde- 
baranium  (Ad)  and  cassiopeium  (Cp).  He  found  their  atomic  weights 
to  be  Ad -172.90  and  Cp  =  174.23. 

The  salts  of  samarium  are  yellow  and  have  a  characteristic  absorp- 
tion spectrum. 

TITANIUM,  ZIRCONIUM,  AND  THORIUM. 

291.  These  uncommon  elements  are  related  to  carbon  and  silicon  in  the 
same  way  as  K,  Rb,  and  Cs  are  to  Li  and  Na,  and  as  Ca,  Sr,  and  Ba 
are  to  Be  and  Mg.  Titanium  and  zirconium  still  give  acid-forming 
oxides,  while  thorium  forms  only  basic  oxides. 

Titanium  displays  very  close  analogy  to  silicon ;  it  frequently  occurs 
with  the  latter,  but  always  in  a  small  amount.  It  is  best  prepared  pure 
by  reducing  TiCl4  with  sodium  in  a  steel  bomb  at  low  red  heat.  Sp.  g.  = 
4.50;  m.-pt.  =  1800°-! 850°.  The  metal  is  hard  and  tough  in  the  cold, 
but  can  be  worked  on  heating;  it  is  a  good  electrical  conductor.  When 
heated  in  a  current  of  nitrogen  it  burns  quantitatively  to  the  nitride  TiN. 
Titanium  dioxide,  TiO2,  occurs  as  mineral  in  three  modifications :  rutile, 
anatase  and  brookite.  Titanium  chloride,  TiCl4,  is  prepared  by  passing 
chlorine  over  the  carbide,  which  is  prepared  in  the  electric  furnace.  TiCl4  is 
liquid  and  fumes  in  the  air  because  of  decomposition  by  atmospheric 
water  into  HC1  and  Ti(OH)4.  Titanic  acid,  Ti(OH)4,  separates  out  as 
a  white  amorphous  powder  when  the  hydrochloric  acid  solution  of  a 
titanate  is  treated  with  ammonia.  This  action  is  due  to  the  weak 
basic  character  of  ammonia  and  the  weak  acid  nature  of  titanic  acid; 
as  a  result  the  ammonium  titanate  is  completely  hydrolyzed  (§  239). 
Like  silicic  and  stannic  acids,  titanic  acid  readily  forms  poly-acids 
(§194).  It  dissolves  in  alkalies  to  form  titanates,  which  are  also 
obtained  by  fusing  TiO2  with  alkalies.  On  the  other  hand  titanic 
acid  dissolves  in  concentrated  sulphuric  acid ;  it  then  remains  in  solution 
even  when  poured  into  water,  because  the  excess  of  sulphuric  acid 
hinders  hydrolytic  dissociation.  Higher  as  well  as  lower  oxides  of 
titanium  are  known.  The  lemon-yellow  oxide  Ti03  is  formed  on  treat- 
ing the  sulphuric  acid  solution  of  Ti(OH)4  with  hydrogen  peroxide  (§  38). 

Zirconium  occurs  in  nature  chiefly  as  zircon,  ZrSiO4.  It  is  not 
reduced  from  the  oxide  by  aluminium.  It  is  prepared  pure  by  heating 
K2ZrF6  withmet  allic  sodium.  Sp.  g.  =6.3.  Small  pieces  burn  bril- 
liantly when  heated  in  the  air.  MOISSAN  obtained  zirconium  carbide, 
CZr,  from  zircon  directly  by  heating  it  with  sugar  charcoal  in  an 


291.]  TITANIUM,  ZIRCONIUM,  AND  THORIUM.  447 

electric  furnace  (1000  amp.  and  40  volts)  for  ten  minutes.  The  silicon 
for  the  most  part  volatilizes.  If  the  carbide  is  treated  with  chlorine 
at  dull-red  heat,  it  is  converted  into  the  chloride.  Zirconium  chloride 
behaves  with  water  in  the  same  way  as  TiCl4  and  SnCl4.  The  hydroxide, 
Zr(OH)4  is  precipitated  by  ammonia  from  acid  solutions  as  a  volu- 
minous mass.  It  is  insoluble  in  alkalies,  but  on  being  fused  with 
the  latter  it  forms  salts  such  as  Na2ZrO3  and  Na4Zr04,  which  are  decom- 
posable by  water.  The  basic  character  of  the  hydroxide  is  apparent 
from  the  fact  that  it  gives  a  sulphate,  Zr(SO4)2,  with  sulphuric  acid, 
which  can  be  recrystallized  out  of  water.  Zirconia,  ZrO2,  emits  a  very 
bright  light  when  heated  strongly. 

Thorium  is  at  present  obtained  mainly  from  monazite  sand;  it  is 
also  found  in  the  thorite  of  Arendal.  It  can  be  prepared  by  electrolysis 
of  a  solution  of  ThCl4  in  molten  alkali  chloride.  It  melts  above  1700°. 
The  hydroxide  Th(OH)4  is  insoluble  in  alkalies.  The  sulphate  crystal- 
lizes with  9H2O, 

Thoria  and  ceria  are  the  essential  constituents  of  the  incandescent 
gas-light  of  A.  VON  WELSBACH.  A  finely  woven  cotton  "  mantle"  is 
saturated  with  a  solution  of  the  nitrates  of  thorium  and  cerium,  in 
which  the  two  are  contained  in  such  a  proportion  that  after  ignition 
the  ash  contains  98-99%  thoria  and  2-1%  ceria.  When  this  ashen 
mantle  is  heated  to  incandescence  by  a  Bunsen  burner  it  gives  out  an 
intense  light.  This  is  apparently  due  to  the  fact  that  such  an  ash. 
mantle  emits  only  a  small  proportion  of  red  rays  and  rays  of  still  greater 
wave-length,  but  mainly  gives  out  rays  of  shorter  wave-length;  hence 
very  little,  if  any,  energy  is  lost  by  the  emission  of  feebly  luminous  rays. 
However,  it  is  found  that  a  mantle  consisting  of  thoria  or  ceria  alone 
or  of  the  two  oxides  in  a  proportion  different  from  the  above  produces 
very  little  light.  So  far  as  the  ceria  is  concerned  this  is  due  to  its  being 
present  in  such  an  excess  that  it  cannot  all  be  raised  by  the  flame  to 
full  incandescence.  An  analogous  phenomenon  is  seen  in  an  ordinary 
flame,  which  when  smoking  (i.e.  when  too  much  carbon  is  present) 
gives  less  light  than  when  not  smoking.  That  it  is  not  the  thoria  which 
emits  the  light  is  proved  by  the  fact  that  a  mantle  consisting  chiefly 
of  ceria  and  containing  only  1-2%  of  thoria  produces  very  little  light. 
We  must  therefore  suppose  that  in  the  mantle  minute  particles  of  ceria 
are  spread  out  upon  the  very  poor  heat-conductor,  thoria;  thus,  since 
their  mass  is  small,  they  are  able  to  reach  the  high  temperature  at  which 
they  emit  the  desired  bright  light ;  for  the  brightness  of  a  flame  increases 
with  about  the  fifth  power  of  the  temperature. 

Thorium  belongs  to  the  radio-active  elements.  When  thorium 
hydroxide  is  dissolved  in  an  acid  and  reprecipitated  with  ammonia,  the 


448  INORGANIC  CHEMISTRY.  [[§§291- 

intensity  of  radiation  is  only  about  45%  of  the  original  intensity.  If 
the  solution  is  evaporated  and  the  ammonium  salts  driven  off  by  ignition, 
a  residue  remains,  too  small  to  weigh;  it  is  thorium  X  and  possesses 
the  remainder  of  the  activity  of  thorium.  The  activity  of  thorium 
increases  again  slowly;  that  of  thorium  X  decreases  and  finally  dis- 
appears. The  half-decay  period  of  the  latter  is  3.64  days.  The  active 
thorium  also  gives  off  an  emanation;  but,  if  in  a  thorium  salt  solution 
the  thorium  is  separated  from  the  thorium  X  (by  repeated  precipitation 
with  ammonia),  the  thorium  is  found  to  have  entirely  lost  its  emanating 
power.  The  thorium  X,  however,  has  strong  emanating  power;  so  we 
conclude  that  the  emanation  is  a  transformation  product  of  thorium  X. 
Between  thorium  X  and  thorium  come  mesothorium  and  radio- 
thorium  as  intermediate  products,  according  to  the  form  of  the  decay 
curve ;  on  the  other  hand,  the  emanation  gives  rise  successively  to  thorium 
A,  B,  and  C. 


VANADIUM,  NIOBIUM  (Columbium),  TANTALUM. 

292.  These  rare  elements  are  allied  to  nitrogen  and  phosphorus  in 
their  properties  and  the  formulae  of  their  compounds.  As  in  all  the  other 
groups  the  metallic  character  becomes  more  prominent  as  the  atomic 
weight  increases.  All  three  are  prepared  by  passing  an  electric  current 
through  rods  'of  thfc  oxides  in  a  vacuum. 

Vanadium  is  characterized  by  an  abundance  of  compound-types. 
There  are,  for  example,  four  chlorides:  VC12,  VCU,  VC14  and  VC1B.  An 
oxychloride,  VOC13,  also  is  known,  which  is  decomposed  by  water  like 
POC13.  The  highest  oxide,  V205,  a  brown  substance,  is  the  starting-point 
for  most  vanadium  preparations.  It  is  an  acid  anhydride,  forming  salts 
which  may  be  derived  from  the  acids  H3V04  (ortho-acid)  and  HV03, 
metavanadic  acid,  and  is  thus  analogous  to  P205. 

Vanadium  occurs  extensively,  though  in  small  quantities.  In  the 
Peruvian  Andes  is  a  large  bed  of  patronite,  a  sulphide,  containing  1 9  per 
cent  of  vanadium.  The  ore  is  roasted  and  fused  with  caustic  soda  and 
saltpetre,  producing  sodium  vanadate,  which  is  extracted  with  water. 
On  saturating  this  solution  with  ammonium  chloride,  ammonium  meta- 
vanadate,  NH4V03,  separates  out  after  a  while  as  a  sandy  powder.  Heat- 
ing converts  it  into  V203.  This  also  serves  as  the  characteristic  test  for 
vanadic  acid. 

Vanadium  melts  at  about  1680°.  It  is  very  active  chemically,  precip- 
itating many  metals  from  their  salt  solutions.  It  is  finding  various  uses, 
both  as  the  brown  oxide  and  in  bronzes  and  vanadium  steels.  It  is 
found  in  meteorites. 


293.]  CHROMIUM  GROUP.  449 

Niobium,  [perhaps  more  justly  called  columbium]  and  tantalum  form 
volatile  chlorides,  NbCl5  and  TaCl5,  which  (like  PC15)  are  decomposed  by 
water.  Particularly  characteristic  of  these  elements  are  their  double 
fluorides,  2KF-NbOF3  and  2KF-TaF5.  The  latter  is  difficultly  soluble, 
the  former  readily  soluble,  in  water.  Use  is  made  of  these  com- 
pounds in  separating  the  two  elements.  The  oxides,  Nb205  and  Ta205 
in  the  presence  of  bases  form  salts  of  niobic  acid,  H3Nb04,  and  tantalic 
acid,  H3TaO4.  The  element  niobium  is  prepared  after  the  method  of 
GOLDSCHMIDT,  by  heating  niobium  pentoxide  with  aluminium  filings. 
The  resulting  product  contains  a  good  deal  of  aluminium,  which  can  be 
volatilized  out  by  heating  in  a  vacuum  to  a  high  temperature.  Pure 
niobium,  obtained  in  this  way,  has  a  specific  gravity  of  12.7  and  melts 
at  about  1950°.  It  is  not  attacked  by  acids  and  burns  only  with  diffi- 
culty in  oxygen. 

Tantalum  is  obtained  by  reducing  its  oxide  with»carbon  in  a  current 
of  hydrogen.  Its  melting-point  is  2850  ±40°. 

CHROMIUM  GROUP. 

Chromium. 

293.  This  element  occurs  principally  in  chromite,  FeO-C^Os 
(§  294),  and  less  commonly  in  crocoite,  PbCrO4.  The  former 
serves  exclusively  for  the  preparation  of  chromium  compounds; 
for  this  purpose  it  is  very  finely  powdered  and  fused  with  an  alkali, 
thus  forming  chromates,  which  are  extracted  with  water. 

The  element  has  been  known  for  a  long  time,  but  it  was  not 
until  1894  that  it  was  prepared  pure  on  a  large  scale  by  MOISSAN. 
He  reduced  chromium  oxide,  Cr2O3,  with  charcoal  in  the  electric 
furnace.  An  easier  method  is  that  of  GOLDSCHMIDT  (§  284),  by 
which  chromium  oxide  is  reduced  with  aluminium  filings.  If  care 
is  taken  to  have  an  excess  of  chromium  oxide  present,  the  metal  is 
obtained  entirely  free  from  aluminium. 

The  metal  thus  obtained  is  lustrous  and  takes  a  polish.  It 
does  not  melt  in  the  oxyhydrogen  flame  but  is  completely  liquefied 
in  the  electric  furnace.  It  boils  at  2200°.  It  does  not  scratch 
glass  (although  the  carbide  C2Cr3  scratches  quartz  and  topaz). 
At  ordinary  temperatures  its  behavior  is  that  of  a  precious  metal, 
i.e.,  it  is  not  in  the  least  affected  by  the  air. 

Chromium  forms  three  sets  of  compounds,  derived  from 
CrO,  chromous  oxide;  Cr2O3,  chromic  oxide;  and  CrOs,  chromic 
anhydride. 


450  INORGANIC  CHEMISTRY.  [§§  293- 

CHROMOUS  COMPOUNDS. 

These  compounds  have  a  very  strong  tendency  to  absorb  oxygen 
and  go  over  into  chromic  compounds;  hence  they  can  only  be  pre- 
served away  from  the  air.  A  solution  of  the  chromous  chloride, 
Crd2,  is  obtained  by  reducing  chromic  chloride,  Cr2Cl6,  with  zinc 
and  sulphuric  acid.  It  has  a  beautiful  blue  color,  which  soon 
turns  to  green  because  of  oxidation.  If  the  solution  of  chromous 
chloride  is  poured  into  a  saturated  solution  of  sodium  acetate, 
chromous  acetate  is  precipitated  as  a  red  crystalline  powder,  which 
is  much  more  permanent  in  the  air  than  the  other  chromium  salts 
and  can  therefore  be  used  for  their  preparation.  The  hydroxide 
Cr(OH)2  is  yellow. 

CHROMIC  COMPOUNDS. 

294.  Chromic  oxide,  Cr2O3,  is  formed  by  heating  chromic 
anhydride,  CrOs,  or  ammonium  chromate  (§  105).  It  can  be 
obtained  crystallized  by  passing  chromyl  chloride,  OO2C12,  through 
a  red-hot  tube.  The  amorphous  compound  is  green;  the  crystals 
are  black.  After  ignition  it  is  insoluble  in  acids.  When  fused 
with  silicates  it  colors  the  latter  green,  whence  its  use  as  a  pigment 
for  coloring  glass  and  china  ware  (chrome  green). 

GUIGNET'S  green,  a  beautiful  pigment,  is  prepared  by  fusing 
potassium  dichromate  (1  part)  with  boric  acid  (3  parts).  The 
potassium  borate  is  dissolved  out  with  water,  leaving  the  coloring 
substance,  Cr203  -2H2O. 

The  hydrogel  of  chromic  oxide,  Cr203-nH2O,  is  precipitated 
when  a  chromium  salt  is  treated  with  ammonia.  It  is  light  blue 
but  dissolves  in  caustic  potash  or  soda  to  a  green  solution.  On 
boiling  this  solution  a  lower  hydrate  of  another  color  is  deposited. 
The  cause  of  this  precipitation  is  readily  understood,  if  it  is  assumed 
that  the  saturated  solution  of  the  lower  hydrate  contains  less 
chromium  ions  than  that  of  the  higher  hydrate.  The  alkaline 
solution  is  thus  supersaturated  in  respect  to  the  lower  hydrate  and 
the  latter  must  be  deposited.  The  solubility  of  chromic  hydroxide 
In  alkalies  shows  its  slightly  acidic  character  £  it  can  also  form  salts 
with  other  metals,  most  of  which  salts  are  derived  from  CrO-OH. 
An  example  of  this  type  is  the  mineral  chromite. 

Chromium  hydroxide  is  only  a  weak  base;    it  does  not  form 


294.]  CHROMIC  COMPOUNDS.  451 

salts  with  weak  acids  such  as  carbonic  acid,  sulphurous  acid,  etc. 
(cf.  §  239). 

Chromium  chloride,  CrCl3,  is  prepared  by  heating  a  mixture 
of  chromic  oxide  and  carbon  in  a  current  of  chlorine;  it  then 
sublimes  in  brilliant  violet  crystal-laminae.  Chromic  chloride  thus 
obtained  dissolves  in  cold  water  very  slowly,  but  if  traces  of 
'chromous  chloride  are  present  it  dissolves  readily.  According  to 
OSTWALD,  this  is  to  be  regarded  as  a'  catalytic  acceleration  of  the 
velocity  of  solution.  The  resulting  solution  is  green;  on  evapora- 
tion green  deliquescent  crystals  of  the  composition  O2Cl6-12H2O 
separate  out.  These  crystals  are  also  obtained  from  the  solution 
of  the  hydrogel  in  hydrochloric  acid.  At  1200-1300°  the  vapor 
density  corresponds  to  the  formula  CrCls. 

Chromic  sulphate,  Cr2(SO4)3,  like  other  chromium  salts 
(nitrate,  chrome  alum,  etc.),  has  the  peculiar  property  of  dissolv- 
ing in  cold  water  to  a  violet  solution,  which  turns  green  on  warm- 
ing. On  cooling,  this  green  color  changes  back  to  violet  (rather 
slowly  with  the  sulphate  solution  but  rapidly  with  other  salts). 
On  slowly  evaporating  the  violet  solution  at  room  temperature  the 
salt  crystallizes  out,  the  sulphate,  for  instance,  with  fifteen  molecules 
of  water;  the  green  solution,  however,  yields  only  an  amorphous 
viscid  mass. 

In  investigating  these  phenomena  the  sulphate  solution  has 
been  usually  employed,  since  its  green  modification  can  be  kept 
the  longest.  Experiments  have  shown  that  the  process  depends 
on  a  splitting  off  of  sulphuric  acid  (1  mol.  H2SO4  from  2  mols. 
sulphate)  and  that  green  "chrom-sulphuric  acids"  are  formed, 
i.e.  substances  with  a  complex  chrom-sulphuric  acid  ion,  since 
they  do  not  give  tests  for  either  chromium  or  sulphuric  acid. 
Thus  only  one-third  of  the  sulphuric  acid  can  be  precipitated  from 
a  green  solution  of  the  sulphate  directly  with  barium  chloride,  or 
in  other  words,  only  a  third  of  the  SO4-ions  of  the  original  violet 
solution  are  still  present.  The  transition  from  the  violet  solution 
to  the  green  can  therefore  be  formulated  in  this  way: 


202(S04)3  +H20  = 

violet  complex  ion 

green 

Only  those  SO4-groups  in  italics  are  precipitated. 


452        •  INORGANIC  CHEMISTRY.  [§§  294- 

In  a  moderate  state  of  dilution  these  chrom-sulphuric  acids  are 
as  strongly  ionized  as  sulphuric  acid  itself. 

An  analogous  behavior  is  shown  by  chromic  chloride,  CrCls. 
A  violet  solution  of  it  can  be  obtained  by  treating  the  violet  solu- 
tion of  the  sulphate  with  the  theoretical  amount  of  barium  chloride. 
The  chlorine  can  then  be  completely  precipitated  with  silver 
nitrate  at  ordinary  temperatures.  If  the  solution  is  boiled  for 
a  time,  however,  and  then  cooled,  silver  nitrate  will  precipitate 
only  two  of  the  three  chlorine  atoms  directly  under  the  same  con- 
ditions; the  third  must  have  gone  with  chromium  to  form  a  com- 
plex ion. 

Chrome  alum,  K2SG4  -Cr2(S04)3-24H2O,  is  best  prepared  by 
passing  sulphur  dioxide  into  a  solution  of  potassium  dichromate 
containing  free  sulphuric  acid: 

K2O2O7 + H2SO4 + 3SO2  =  K2S04  •  Cr2(SO4)3  +  H20. 

It  can  be  obtained  in  finely  developed  octahedrons  with  edges 
several  centimeters  in  length.     It  is  used  in  dyeing  and  tanning. 

CHROMA  TES. 

295.  There  are  numerous  salts  of  the  formula  M2Cr04  which 
are  derived  from  the  oxide  CrOs,  chromic  anhydride,  as  in  the 
similar  case  of  sulphuric  anhydride,  SOs;  but  while  in  the  latter 
case  the  sulphuric  acid  itself,  H2SO4,  is  also  a  stable  compound, 
chromic  acid,  H2CrO4,  has  not  as  yet  been  isolated.  When  an 
acid  is  added  to  a  chromate  only  the  anhydride  is  obtained;  the 
acid  H2Cr04  breaks  up  forthwith  into  water  and  anhydride.  The 
salts  of  chromic  acid  are  isomorphous  with  the  corresponding 
sulphates. 

Chromic  anhydride  is  obtained  on  adding  sulphuric  acid  to  a 
concentrated  solution  of  potassium  dichromate;  it  separates  out 
in  the  form  of  long  red  rhombic  needles,  which,  when  freed  from 
all  sulphuric  acid,  do  not  deliquesce  in  the  air.  They  are  readily 
soluble  in  water.  Heating  to  250°  breaks  them  up  into  chromic 
oxide  and  oxygen: 

2Cr03=Cr2O3  +  30. 

Chromium  trioxide  is  a  very  powerful  oxidizing-agent;  its  solu- 
tion cannot  be  filtered  through  paper  because  it  destroys  the  paper 


295.]  CHROMATES.  453 

by  oxidation.  When  strong  alcohol  is  dropped  on  chromic  anhy- 
dride it  takes  fire,  chromic  oxide  (Cr2O3)  being  formed  at  the  same 
time.  Hydrochloric  acid  is  oxidized  to  chlorine  and  water,  sulphur- 
ous acid  to  sulphuric  acid.  On  passing  dry  ammonia  over  CrO3 
crystals,  the  gas  takes  fire  and  the  oxide  is  reduced.  When  heated 
with  sulphuric  acid  it  yields  oxygen  and  chromium  sulphate. 
Hydrogen  sulphide  reduces  the  aqueous  solution,  sulphur  being 
deposited.  Chromic  anhydride  thus  displays  various  characteristics 
of  peroxides  such  as  Pb02,  BaO2,  etc. 

In  addition  to  the  normal  salts  of  chromic  acid,  e.g.  K2CrO4, 
there  are  also  d  i  chromates,  t  r  i  chromates,  etc.,  which  can  be 
regarded  as  combinations  of  one  molecule  of  the  neutral  salt  with 
one  or  more  Cr03  molecules  : 

K2CrO4  +  CrO3  =  K2Cr207  ;  K2Cr  O4  +  2Cr03  =  K2Cr3Oi0,  etc. 

Dichromate  Trichromate 

If  J  molecule  of  sulphuric  acid  is  added  to  1  molecule  of 
chromate,  the  yellow  color  of  the  chromate  solution  is  changed  to 
the  red  color  of  the  dichromate;  a  CrO4-ion  gives  up  its  electrical 
charge  and  an  atom  of  oxygen  to  the  hydrogen  ions  of  the  free 
acid,  thus  yielding  water  and  forming,  together  with  a  second 
Cr04-ion,  the  red  ion  O2O7: 

4K'  +  2CrO4"  +  2H"  +J§04"  =  4K*  +  Cr207"  +  H20  +  S04". 


Acid  salts  of  chromic  acid  do  not  exist  on  account  of  this  reducing 
effect  of  the  hydrogen  ions  on  the  OO4-ions,  for  which  reason  also 
free  chromic  acid,  H2CrO4,  is  incapable  of  independent  existence. 

Chromic  acid  is  a  weak  acid,  since  its  insoluble  (in  water)  salts, 
e.g.  those  of  barium,  lead  and  silver,  are  readily  dissolved  by 
strong  acids  (§  146). 

Alkali  chromates  are  invariably  obtained  by  fusing  a  chromium 
compound  with  an  alkali  carbonate  and  an  oxidizing-agent.  The 
latter  is  unnecessary  when  the  fused  mass  can  be  brought  suffi- 
ciently in  contact  with  the  oxygen  of  the  air  by  stirring.  Chromite 
is  worked  up  commercially  into  chromates  in  this  way;  it  is  cal- 
cined with  soda  and  lime  above  1000°  in  a  reverberatory  furnace: 

2Cr2O3  -  FeO  +  4Na2CO3  +  4CaO  +  7O 

=  4Na2Cr04  +4CaC03  +  Fe203. 


454  INORGANIC  CHEMISTRY.  [§§295- 

The  resulting  sodium  chromate  is  lixiviated  and  sulphuric  acid 
is  added  to  its  solution;  on1  evaporation  sodium  dichromate, 
Na2Cr2O7,  crystallizes  out,  and  this  can  be  converted  into  potassium 
dichromate,  a  well-known  salt,  by  double  decomposition  with 
potassium  chloride. 

The  fusion  is  much  more  readily  accomplished  when  caustic 
potash  (KOH)  is  used  instead  of  soda  (Na2CO3),  probably  because 
fused  potassium  hydroxide  absorbs  oxygen  from  the  air  and  forms 
the  peroxide,  thus  becoming  a  much  more  active  oxygen-carrier 
than  soda.  Under  these  circumstances  the  oxidation  proceeds 
rapidly  and  completely  as  low  as  500°. 

Potassium  dichromate  finds  frequent  use  as  an  oxidizing-agent 
in  sulphuric  acid  solution,  being  itself  reduced  to  chromic  sulphate: 

K2Cr207 + 4H2SO4  =  K2S04 + O2  (S04)3  +  4H20 + 3O. 

An  important  commercial  task  is  the  regeneration  of  the  chromic  acid 
from  such  a  chromium  sulphate  solution.  The  method  followed  in  the 
factories  at  Hochst,  Germany,  is  an  electrical  one.  The  solution  is  elec- 
trolyzed  between  lead  electrodes  in  a  vessel  containing  a  diaphragm 
(porous  partition).  By  the  action  of  the  current  chromic  acid  is  formed 
at  the  anode,  while  at  the  cathode  hydrogen  is  evolved.  Besides  this,  a 
change  occurs  in  the  concentration  of  the  sulphuric  acid  on  both  sides  of 
the  diaphragm;  it  becomes  higher  in  the  anode  portion,  and  lower  in 
the  cathode  portion.  The  liquid  oxidized  at  the  anode  can  be  used  for 
oxidizing  purposes  without  any  further  preparation.  The  chromic  acid 
is  again  reduced  to -chromic  oxide,  Cr^,  and  the  reduced  liquid  is  then 
introduced  into  the  cathode  portion,  while  the  liquid  which  previously 
occupied  that  space,  is  brought  over  to  the  anode  side  of  the  diaphragm. 
When  the  current  is  again  turned  on,  the  liquid  at  the  cathode,  which 
at  the  beginning  of  this  second  operation  is  richer  in  sulphuric  acid  than 
the  liquid  at  the  anode,  yields  its  surplus  to  the  cathode  liquid.  In  this 
way  an  accumulation  of  sulphuric  acid  is  avoided,  and  the  same  liquid 
can  really  be  used  continuously  as  an  oxidizing-agent. 

The  chromates  are  yellow  (except  silver  chromate,  which  is  red) 
and  the  dichromates  are  red.  Lead  chromate,  PbOO4,  is  insoluble 
in  water  and  is  used  as  a  pigment  (chrome  yellow). 

On  heating  potassium  dichromate  with  potassium  chloride  and  sul- 
phuric acid  a  dark-red  liquid  distils  over,  which  has  the  composition 
OO2C12  and  the  boiling-point  117°  and  must  be  considered  as  the  chloride 


296.]  MOLYBDENUM.  455 

of  chromic  acid;  it  is  called  chromyl  chloride,  or  chromium  oxy chloride: 
K2Cr207 +4KC1 +3H2S04  =2CrO2Cl2 +3K2S04 +3H20. 

Water  breaks  it  up  into  CrO3  and  HC1. 

Cl 

The  semi-chloride  of  chromic  acid,  O02<       ,  is  known  only  in  the 

form  of  salts.    The  potassium  salt,  for  example,  is  obtained  by  heating 
potassium  dichromate  with  concentrated  hydrochloric  acid: 

K20207  +2HC1  =2Cr02<^K-fH20. 

It  crystallizes  in  red  prisms. 

On  treating  a  chromic  acid  solution  with  hydrogen  peroxide  in  excess 
a  beautiful  blue  coloration  appears,  which  is  absorbed  by  ether  on  shak- 
ing. It  is  due  to  a  perchromic  acid,  whose  ammonium  salt  NH4OO5-  H202 
can  be  isolated  as  a  violet-black  powder  similar  to  powdered  potassium 
permanganate.  In  concentrated  aqueous  solution  decomposition  soon 
occurs,  the  dichromate  being  formed  and  oxygen  given  off. 

Molybdenum. 

296.  This  comparatively  rare  element  is  found  in  nature  in 
molybdenite,  MoS2,  and  wulfenite,  PbMoC>4.  The  former  alone  is 
used  in  preparing  molybdenum  and  its  compounds.  It  is  roasted 
and  so  converted  into  the  trioxide,  MoOs. 

The  element  itself  is  obtained  from  its  oxides  or  chlorides 
by  heating  them  red-hot  in  a  current  of  hydrogen.  The  product 
is  a  steel-gray  powder  which  fuses  with  great  difficulty  to  a  silvery 
metallic  mass.  Sp.  g.  =8.6.  Heating  in  the  air  converts  it  into 
the  trioxide.  It  is  not  attacked  by  hydrochloric  or  dilute  sul- 
phuric acid,  but  is  readily  dissolved  by  nitric  and  concentrated 
sulphuric  acid.  Molybdenum  also  has  recently  found  a  metallur- 
gical use  in  varying  the  properties  of  steel. 

This  element  is  noted  for  the  great  variety  of  its  compounds; 
some  of  the  more  important  ones  may  be  mentioned  here. 

In  addition  to  the  oxides  Mo2O3  (weakly  basic)  and  Mo02  (in- 
different) there  is  molybdenum  trioxide,  MoOa,  which,  like  CrOs, 
is  an  acid  anhydride.  It  is  a  white  powder  which  turns  yellow  on 
heating.  It  is  very  sparingly  soluble  in  water.  With  alkalies  it 
forms  molybdates.  It  has  a  tendency  to  form  p  o  I  y-molybdates 
even  stronger  than  the  similar  tendency  of  chromic  anhydride; 


456  INORGANIC  CHEMISTRY.  [§§  296- 

ammonium  heptamolybdate,  (NH4)6Mo7024-4H20  (derivable  from 
the  acid  7H2MoC>4  —  4H2O),  commonly  known  as  "ammonium 
molybdate,"  is  a  typical  example.  The  addition  of  a  strong  acid 
to  a  molybdate  solution  precipitates  white,  glistening  crystal^lam- 
inae  of  molybdic  acid,  H2MoO4,  which  dissolve  in  an  excess  of  acid. 
A  solution  thus  prepared  from  ammonium  molybdate  and  an  excess 
of  nitric  acid  serves  as  a  test-reagent  for  phosphoric  acid,  with 
which  it  forms  a  yellow  precipitate  of  about  the  composition 
(NH4)3PO4-14MoO3  +  4H2O  on  warming  (cf.  §§  146  and  162). 

Of  the  chlorides  the  compounds  Mods,  MoCU,  and  MoCl5  are 
known.  In  the  oxychlorides  MoOCU  and  MoO2Cl2  molybdenum 
can  be  regarded  as  sexivalent. 

The  chloride  MoCl2  does  not  exist  according  to  MUTHMANN  (neither 
does  MoO) ,  but  a  chloride  Mo3Cl6  is  known. 

A  very  characteristic  test  for  molybdic  acid  (the  most  common 
molybdenum  compound)  is  the  following:  the  substance  is  mixed 
with  zinc  and  sulphuric  acid ;  at  first  a  blue  coloration  (a  molybdate 
of  molybdic  oxide)  appears  but  it  soon  turns  green  and  then  brown. 
This  brown  coloration  is  due  to  a  salt  of  the  oxide  Mo203. 

TUNGSTEN. 

297.  The  minerals  in  which  this  element  chiefly  occurs  are  scheelite, 
CaWO4,  wolframite,  or  wolfram,  (Fe,  Mn)WO4,  and  hubnerite,  MnW04. 
The  metal  is  obtained  pure  by  the  method  of  GOLDSCHMIDT  (§  284),  i.e., 
by  the  reduction  of  pure  tungstic  acid  with  aluminium  filings.  The 
metal  so  obtained  is  very  pure;  sp.  g.,  18.73;  melting-point  somewhat 
above  2800°.  It  is  malleable  and  scratches  glass.  In  combination  with, 
carbon  it  is  very  much  harder.  It  is  very  permanent  in  the  air.  Sul- 
phuric acid,  hydrochloric  acid,  aqua  regia,  and  hydrofluoric  acid  attack 
it  very  slowly,  but  it  rapidly  dissolves  in  a  mixture  of  hydrofluoric  and 
nitric  acids.  Fused  caustic  potash  dissolves  it  slowly  with  the  evolution 
of  hydrogen.  Tungsten  is  employed  in  the  iron  industry,  since  a  small 
percentage  of  tungsten  increases  the  hardness  of  steel  in  a  marked  degree 
(tungsten,  or  wolfram,  steel) .  Extremely  fine  wires  of  the  metal  are 
made  use  of  in  some  of  the  newer  incandescent  electric  lights. 

Tungsten,  like  chromium  and  molybdenum,  is  also  characterized  by 
an  abundance  of  compound-types.  The  chlorides  WC12,  WC14,  WC15, 
and  WC16  are  known  to  exist.  The  lower  ones  are  prepared  from  the 


298.]  URANIUM.  457 

hexachloride  by  heating  in  a  current  of  hydrogen  or  carbon  dioxide.  The 
hexachloride  itself  is  formed  by  direct  synthesis;  it  is  a  violet-black 
crystalline  substance;  water  converts  it  into  the  anhydride,  WO3. 

Tufogstic  anhydride,  WO3,  is  obtained  by  precipitating  the  hot  solution 
of  a  tungstate  with  nitric  acid.  It  is  insoluble  in  water  and  acids  but 
soluble  in  alkalies.  The  addition  of  an  acid  to  the  cold  solution  of  a 
tungstate  precipitates  tungstic  acid,  WO(OH)4[=W(OH)6-H20].  The 
latter  forms  polyacids,  like  chromic  and  molybdic  acids.  Like  molybdic 
acid  also  it  has  the  property  of  uniting  with  phosphoric  and  arsenic  acids 
to  form  complex  phospho-tungstates  and  arseni-tung- 
states  . — The  following  is  a  very  characteristic  test  for  tungstates:  If 
stannous  chloride  is  added  to  a  tungstate  solution,  a  yellow  precipitate 
(WO3)  is  produced.  On  the  addition  of  hydrochloric  acid  and  warming, 
a  beautiful  blue  solution  (W205)  is  obtained. 

URANIUM. 

298.  The  principal  uranium  mineral  is  uraninite,  which  usually  con- 
tains some  iron.  The  m  e  t  a  1  is  obtained  by  heating  the  chloride  with 
sodium  or  by  the  electrolysis  of  the  chloride  or  by  the  reduction  of  the 
oxide  with  carbon  in  the  electric  furnace.  It  is  silvery-white  and  has  a 
specific  gravity  of  18.7.  It  is  much  more  volatile  than  iron  in  the  electric 
furnace.  When  it  is  in  the  form  of  a  fine  powder  it  burns  in  a  current  of 
oxygen  as  low  as  170°.  In  the  same  state  it  decomposes  water  slowly 
at  room  temperature.  When  nitrogen  is  passed  over  uranium  the  two 
elements  combine  readily  at  1000°  to  form  a  yellow  nitride.  Another 
interesting  compound  is  the  carbide  C3U2  (obtained  from  uranium  oxide 
and  charcoal  in  the  electric  furnace),  inasmuch  as  the  addition  of  water 
yields  not  only  methane  but  liquid  and  solid  hydrocarbons. 

Uranium  forms  two  sets  of  compounds;  in  the  ous  compounds  it  is 
quadrivalent  (UX4),  in  the  ic  compounds  sexivalent  (UX6).  The  former 
pass  readily  into  the  latter.  The  oxide  UO2  has  an  exclusively  basic  char- 
acter; it  is  obtained  by  igniting  the  other  oxides  in  a  current  of  hydro- 
gen. It  was  at  one  time  regarded  as  the  metal  itself. 

Uranic  oxide,  UO3,  is  a  yellow  powder,  prepared  by  heating  the 
nitrate.  The  corresponding  hydroxide,  U(OH)6,  is  not  known,  but  salts 
of  the  compound  U(OH)6-2H20=UO2(OH)2  with  acids  have  been 
prepared.  Since  the  U02  group  acts  here  as  a  bivalent  radical  it  is 
called  uranyl  and  its  salts  uranyl  salts,  e.g.  U02(N03)2,  uranyl  nitrate, 
crystallizing  with  6H2O  in  beautiful  greenish-yellow  prisms.  Uranium 
trioxide  also  has  somewhat  the  character  of  an  acid  anhydride;  if  caustic 
potash  and  soda  are  added  to  uranyl  salt  solutions  yellow  uranates 


458 


INORGANIC  CHEMISTRY. 


[§§  298- 


(K2U2O7  and  Na^O?)  are  precipitated,  which  are  soluble  in  acids. 
Uraninite  can  be  regarded  as  the  uranate  of  uranous  oxide,  U3O8 
=2UO3-UO2.  Both  oxides  are  converted  into  this  U3O8  oxide  by  heat- 
ing in  the  air.  Uranium  salts  are  used  to  impart  to  glass  a  beautiful 
greenish-yellow  fluorescence. 

The  detection  of  uranyl  salts  is  accomplished  with  the  aid  of  the  pre- 
cipitate, soluble  in  excess,  which  they  give  with  ammonium  carbonate 
and  by  the  reddish-brown  precipitate  with  potassium  ferrocyanide.  For 
the  radioactive  properties  of  uranium  see  §  267. 

SUMMARY  OF  THE  GROUP. 

299.  The  elements  chromium,  molybdenum,  tungsten,  and  ura- 
nium, in  connection  with  sulphur,  constitute  a  natural  group  in  the 
periodic  system.  Particularly  in  the  higher  oxides  there  is  con- 
siderable analogy  with  the  behavior  of  this  metalloid.  Their  acids, 
for  example,  all  have  the  formula  H2RO4.  Moreover  sulphur  also 
has  the  ability  to  form  polyacids  (pyrosulphuric  acid)  although  it 
is  not  so  prominent  as  in  the  first-named  four  elements.  Several 
of  their  salts  are  isomorphous.  The  strength  of  the  acids  decreases, 
as  in  other  groups,  with  rising  atomic  weight.  Another  character- 
istic of  all  the  elements  of  this  group  is  the  great  abundance  of 
formula  types;  it  is  also  very  noticeable  in  the  case  of  sulphur, 
whose  acids  are  remarkably  numerous.  The  physical  properties  of 
these  elements  have  not  yet  been  fully  determined,  but  a  few  of 
them  are  given  in  the  following  table: 


Cr 

Mo 

w 

U 

Atomic  weight 

52  0 

96  0 

184  0 

238  5 

Specific  gravity          .  . 

6  7 

8  6 

16  6 

18  7 

Color                

white 

white 

white 

white 

Melting-point          

>  2800° 

MANGANESE. 

300.  This  element  is  widely  diffused  in  nature.  Its  most  im- 
portant minerals  are  pyrolusite,  MnC>2,  hausmannite,  MnsO^  and 
rhodochrosite,  MnCOs. 

The  metal  is  of  minor  importance.  It  is  best  prepared 
by  the  GOLDSCHMIDT  method,  i.e.  by  reducing  pyrolusite  with 
aluminium  powder,  when  it  is  obtained  as  a  regulus  of  brilliant 


300.]  MANGANESE.  459 

lustre.  Sp.  g.  =  7.2-8.0;  m.-pt.=  1245°;  b.-pt.=  1900°.  It  under- 
goes surface  oxidation  readily  in  moist  air,  which  gives  the 
regulus  an  iridescence,  and  whea  finely  divided  decomposes 
boiling  water.  It  dissolves  in  acids  to  form  manganous  salts. 

Manganese  forms  several  series  of  compounds:  the  manganov* 
compounds  of  the  type  MnX2;  the  manganic  compounds,  MnX3; 
manganic  acid,  H2MnO4,  which  can  be  derived  from  an  anhydride 
MnOs ;  permanganic  acid,  HMnO4  derivable  from  the  oxide  Mn2O7. 
Most  of  the  familiar  salts  of  this  element  are  derived  from  man- 
ganous oxide,  MnO.  This  oxide,  which  is  prepared  by  heating  the 
carbonate  in  the  absence  of  air,  is  an  amorphous  green  powder,  that 
oxidizes  readily  in  the  air  to  the  higher  oxide  Mn304.  Manganous 
hydroxide,  Mn(OH)2,  is  white  when  freshly  precipitated  from  solu- 
tions by  an  alkali  but  soon  turns  brown  in  the  air  because  of  the 
formation  of  manganic  hydroxide,  Mn2(OH)6- 

The  solutions  of  manganous  salts  are  pink  (color  of  the  Mn*»- 
ion).  The  chloride,  MnCl2,  crystallizes  with  four  molecules  of 
water.  It  can  be  obtained  anhydrous  by  heating  the  double  salt 
MnCl2-2NH4Cl  +  H2O,  since  the  hydrochloric  acid  set  free  hinders 
the  hydrolytic  dissociation  of  the  chloride.  The  sulphate,  MnSC>4, 
crystallizes  below  6°  with  7H20,  above  this  temperature  with 
5H2O.  It  forms  double  salts,  such  as  K2SO4-MnS04+6H20,  simi- 
lar to  those  of  magnesium  and  iron;  they  are  moreover  isomor- 
phous  with  the  latter. 

Manganous  sulphide,  MnS,  has  a  pinkish-white  color,  which 
distinguishes  it  from  all  other  sulphides. 

If  ammonium  chloride  is  added  to  the  solution  of  a  manganese 
salt,  no  hydroxide  is  precipitated  by  ammonia;  this  is  analogous 
to  what  is  observed  with  magnesium  (§  254).  The  solution  is, 
however,  readily  oxidized  by  the  oxygen  of  the  air  and  brown 
manganic  hydroxide  is  deposited. 

The  manganic  ion  Mn"*  is  only  weakly  basic.  Its  salts  are 
almost  completely  hydrolyzed  in  aqueous  solution.  The  sulphate 
gives  alums  with  cesium  and  rubidium  sulphates,  which  are  also 
very  unstable. 

Manganic  oxide,  Mn2C>3,  is  obtained  from  any  of  the  other 
oxides  by  heating  in  an  oxygen  current.  Since  dilute  sulphuric 
acid  reacts  with  it,  giving  manganous  sulphate  and  manganese 
dioxide,  the  oxide  Mn2C>3  is  often  considered  as  MnO-Mn02.  The 


460  INORGANIC   CHEMISTRY.  [§§  300- 

corresponding  hydroxide  is  soluble  in  cold  hydrochloric  acid  to  a 
dark-brown  solution.  It  is  not  certain  whether  this  solution  con- 
tains Mn2Cl6  or  MnCl2  and  MnCl4;  on  being  warmed  it  gives  off 
chlorine  and  is  then  known  to  contain  the  manganous  chloride. 

Mangano-manganic  oxide,  Mn304  or  MnO-Mn2O3,  is  obtained 
on  strongly  igniting  the  other  oxides  in  the  air.  It  is  a  brownish- 
red  powder.  When  heated  with  hydrochloric  acid  it  yields  chlorine. 

Manganese  di-  (or  per-)  oxide,  MnO2,  the  best- known  man- 
ganese mineral  (pyrolusite) ,  is  commercially  of  great  importance 
in  the  production  of  chlorine.  In  the  cold  it  dissolves  in  hydro- 
chloric acid  to  a  very  dark  liquid,  probably  containing  the  tetra- 
chloride,  and  gives  off  no  chlorine;  when  warmed  it  decomposes 
into  chlorine  and  manganous  chloride  (§  25). 

Since  pyrolusite  is  comparatively  expensive;  various  methods 
have  been  devised  for  reconverting  the  manganous  chloride  into 
the  peroxide.  One  which  is  of  practical  importance  is  the  WELDON 
process.  An  excess  of  milk  of  lime  is  added  to  the  chloride  solution, 
whereupon  air  is  forced  through  the  warmed  liquid.  'The  manganous 
hydroxide  which  is  precipitated  undergoes  oxidation  and  is  con- 
verted into  calcium  manganite,  CaMn03  (  =  CaO-MnO2),  which 
settles  down  as  a  black  slimy  mass: 

MnCl2 + 2CaO  +  0  =  CaMn03  +  CaCl2. 

The  calcium  chloride  solution  is  run  off  and  the  manganite  is  used 
for  generating  chlorine,  since  it  acts  towards  hydrochloric  acid  like 
a  mixture  of  lime  and  manganese  dioxide. 

The  value  of  the  peroxide  depends  on  the  amount  of  chlorine 
it  can  produce  with  hydrochloric  acid.  In  order  to  determine  this, 
the  mineral,  finely  pulverized,  is  warmed  with  hydrochloric  acid 
and  the  evolved  chlorine  passed  into  potassium  iodide  solution, 
whereupon  an  equivalent  amount  of  iodine  is  liberated.  This 
iodine  can  be  titrated  with  thiosulphate  (§  93). 

Manganic  acid  and  Permanganic  acid* 

301.  When  manganese  compounds  are  fused  with  potassium 
hydroxide  in  the  air  or,  better,  in  the  presence  of  an  oxidizing-agent 
(potassium  nitrate  or  chlorate)  a  green  mass  results,  which  is  dis- 
solved by  cold  water,  forming  a  dark-green  solution.  On  evaporat- 
ing this  solution  in  a  vacuum  dark-green  rhombic  prisms  of  potas- 


301.]  MANGANIC  AND  PERMANGANIC  ACIDS.  461 

slum  manganate,  K2Mn04,  crystallize  out,  which  have  a  metallic 
lustre  and  are  isomorphous  with  potassium  chromate.  They 
dissolve  in  potassium  or  sodium  hydroxide  solutions  without 
change,  but  are  decomposed  by  water  with  the  separation  of  man- 
ganese dioxide  and  the  formation  of  potassium  permanganate, 
KMn04,  the  latter  giving  the  solution  a  deep  violet  color: 

3K2MnO4  +  3HoO  =  2KMn04 + Mn02 .  H20 + 4KOH. 

On  account  of  these  changes  of  color  the  manganate  solution 
received  the  name  chamceleon  minerale,  from  the  early  chemists. 

Both  in  the  solution  of  a  manganate  and  in  that  of  a  perman- 
ganate we  have  the  anion  Mn04;  in  the  former,  however,  it  is 
bivalent,  in  the  latter  univalent.  This  causes  the  difference  in 
the  properties  of  the  two  ions ;  the  univalent  ion  Mn04'  is  deep  red 
and  resembles  the  perchloric  acid  ion  in  behavior,  while  the  bivalent 
MnO4"  is  deep  green  and  displays  analogy  to  the  SO4"  ion  of 
sulphuric  acid.  The  bivalent  ion  MnO4"  is  only  stable  in  alkaline 
liquids;  it  is  converted  by  water  (more  easily  by  acids)  into  the 
univalent  ion: 

3K2Mn04  +  4HN03  -  2KMn04  +Mn02  +  4KN03 + 2H20, 
or,  written  in  ions: 

6K*  +3MnO4" + 4H" + 4NO3'  -  2K" 

+2MnO4'+Mn02+4K'  +  4N03'+2H20. 

The  reaction  obviously  amounts  to  a  formation  of  water  by  the 
four  hydrogen  ions  and  two  oxygen  atoms  which  they  extract  from 
a  bivalent  anion  Mn04",  the  latter  being  reduced  to  Mn02.  Of 
the  four  negative  charges  which  are  required  to  neutralize  the  four 
positive  charges  of  the  hydrogen  ions  two  are  taken  from  this  MnO4 
anion,  which  is  reduced  to  Mn02,  and  the  remaining  two  from  two 
other  bivalent  anions  Mn04",  which  thus  become  univalent.  The 
transformation  of  potassium  manganate  into  the  permanganate  is 
effected  commercially  by  passing  ozone  into  its  concentrated  solution: 

2K2Mn04+03  =2KMnO4+K20  +  02. 

The  permanganate  crystallizes  out  of  the  solution  and  the  resulting 
mother-liquor  can  at  once  be  used  with  a  fresh  quantity  of  pyrolusite 
to  prepare  more  manganate. 

Potassium  permanganate,  KMn04,  crystallizes  in  beautiful 
glistening  greenish-black  prisms  of  the  rhombic  system,  which 


462  INORGANIC    CHEMISTRY.  [§§  301- 

dissolve  readily  in  water,  forming  a  deep-violet  liquid.  This  salt 
is  isomorphous  with  potassium  perchlorate.  All  solutions  of  per- 
manganates display  the  same  absorption  spectrum,  viz.,  five  dark 
bands  in  the  yellow  and  green,  no  matter  what  the  base  is.  It  is 
thus  evident  that  the  ion  Mn04'  is  really  the  coloring-agent.' 

The  solution  of  potassium  permanganate  acts  as  a  powerful 
oxidizing-agent;  in  acid  solutions  two  KMnO4  molecules  yield 
five  oxygen  atoms  : 

2KMn04  +  3H2S04  =  K2S04  +  2MnSO4  +  3H20  +  50. 

The  process  may  be  regarded  as  a  transformation  of  the  anhydride 
of  permanganic  acid,  Mn207(=2HMnO4—  H2O),  into  two  mole- 
cules of  basic  oxide,  MnO,  and  five  atoms  of  oxygen;  thus: 


In  neutral  or  alkaline  solutions,  however,  two  KMn04  molecules 
yield  only  three  atoms  of  oxygen,  manganese  peroxide  being 
deposited  at  the  same  time  (transformation  of  Mn2O7  into 
2Mn02+30): 

2KMnO4  +  H20  =  2MnO2  +  2KOH  +  30. 

Since  in  oxidations  with  potassium  permanganate  in  acid  solu- 
tion the  deep  color  of  the  permanganate  is  replaced  by  the  very 
faint  color  of  manganous  sulphate,  many  substances  can  be  titrated 
with  potassium  permanganate  in  acid  solution  without  an  indi- 
cator. Ferrous  sulphate  is  oxidized  to  ferric  sulphate;  oxalic 
acid  goes  over  into  carbon  dioxide  and  water;  nitrous  acid  in  very 
dilute  solutions  is  converted  into  nitric  acid  (§  126)  ;  from  hydrogen 
peroxide  water  and  oxygen  gas  are  formed.  All  these  reactions 
proceed  quickly  and  quantitatively  at  ordinary  temperatures  so 
that  they  are  suitable  for  titration. 

Permanganic  acid  is  known  only  in  aqueous  solution;  however,  its 
anhydride,  Mn2O7,  can  be  obtained.  It  is  prepared  by  carefully  treat- 
ing dry  permanganate  with  concentrated  sulphuric  acid.  It  is  a  vola- 
tile, brownish-green,  oily  liquid,  whose  vapor  explodes  easily,  yielding 
oxygen  and  manganese  dioxide. 

Manganese  occupies  an  isolated  position  in  the  periodic  system. 
No  elements  are  known  which  are  related  to  it  as  the  elements 
Mo,  W  and  U  are  to  chromium.  Moreover,  only  in  its  highest 
stage  of  oxidation,  permanganic  acid,  does  it  display  analogy  with 


302.]  IRON.  463 

the  corresponding  chlorine  compound,  HC104.    The  salts  of  both 
acids  are  isomorphous  and  both  are  powerful  oxidizing-agents. 

IRON. 

302.  Iron  is  the  most  useful  metal,  and  is  therefore  prepared 
commercially  on  an  enormous  scale  (approximately  50,000,000 
metric  tons  a  year).  It  occurs  only  rarely  native ,  e.g.  in  meteoric 
rocks.  In  the  form  of  oxides,  sulphides  and  silicates  it  is  widely 
diffused  in  nature  and  is  found  in  very  large  quantities.  The 
most  important  minerals  for  the  iron  industry  are  magnetite T 
FeaQ4,  hematite,  FeaOs,  and  .s?Vfcr?!feT  Ff>f!O3.  The  pyrites  (FeS2; 
etc.)  are  worked  up  into  iron  after  they  have  been  roasted  in  the 
sulphuric  acid  factories. 

The  metallurgy  of  iron  is  theoretically  very  simple;  it  is  based 
on  the  ability  of  carbon  to  reduce  the  oxides  of  iron  to  the  metal 
at  an  elevated  temperature.  This  process  (smelting)  is  car- 
ried out  in  blast  furnaces. 

The  iron  ore  is  first  roasted  (calcined)  to  remove  volatile 
substances  (H2O,  C02,  S,  As,  etc.)  and  loosen  up  the  mineral. 
Then  it  is  crushed  and  mixed  with  a  slag-forming  substance 
(flux,  see  §  242),  according  to  the  grade  of  the  ore.  If  the  gangue, 
or  earthy  matrix,  contains  much  silica  or  alumina,  limestone  or 
dolomite  is  employed  as  the  fluxing-agent,  but  ores  rich  in  lime 
or  magnesia  are  mixed  with  quartz  or  aluminous  ore  to  effect  the 
necessary  fusion  and  formation  of  slag  (silicates  of  Al,  Mg  and  Ca). 

The  blast  furnace,  previously  warmed  to  the  proper  tempera- 
ture or  already  in  operation,  is  charged  from  above  with  alternate 
layers  of  coke  and  the  mixture  of  ore  and  flux,  both  being  intro- 
duced in  "rounds/ '  or  "charges,"  of  definite  weight.  (Sometimes 
charcoal  or  anthracite  is  used  as  fuel.)  The  modern  furnaces  (Fig. 
73)  are  built  of  fire-brick  encased  in  iron  and  are  of  much  lighter 
construction  than  those  formerly  used.  They  vary  greatly  in  size 
but  consist  mainly  of  a  long  shaft  tapering  towards  both  ends. 
In  order  to  utilize  the  escaping  hot  gases  (CO,  etc.)  an  apparatus 
("cup  and  cone")  is  fitted  on  the  top  to  conduct  them  off  and 
also  allow  the  introduction  of  the  charge.  The  air  necessary  for 
the  process  is  forced  in,  hot,  through  pipes  (twyers)  at  the  bottom 


464  INORGANIC  CHEMISTRY.  [§§  302- 

of  the  furnace.    The  burning  coke  produces   carbon  monoxide, 
which  is  the  principal  factor  in  the  reduction  of  the  ore: 

Fe2O3  +  SCO  =  2Fe  +  3C02. 

The  reduced  iron  sinks  downward  and  comes  in  contact  with 
carbon  at  a  high  temperature;  as  a  result  some  of  the  carbon  is 


FIG.  73. — BLAST  FURNACE. 

dissolved  by  it  and  its  melting-point  considerably  depressed. 
When  a  definite  stage  is  reached  the  fused  iron  is  drawn  off  below. 
It  is  protected  from  atmospheric  oxidation  by  the  slag  floating 
on  it. 


303.]  IRON,  465 

The  attempt  to  extract  iron  from  its  ores  by  electrical  heating 
has  met  with  success.  STASSANO  calculates  on  the  basis  of  the 
analysis  of  the  ore  the  additions  which  will  be  necessary  to  yield 
a  slag  as  nearly  as  possible  of  the  composition  Si02  +  4  Base,  com- 
presses the  finely  powdered  material  to  briquetts  with  the  aid  of 
tar  in  hydraulic  presses,  and  smelts  it  in  a  specially  constructed 
arc  furnace. 

303.  It  was  stated  above  that  the  waste  furnace-gases  contain  a  con- 
siderable quantity  of  carbon  monoxide ;  therefore  a  large  amount  of  heat 
is  lost,  which  could  be  utilized  by  burning  the  monoxide  to  dioxide. 
Supposing  that  this  incomplete  reaction  was  due  to  an  incomplete  con- 
tact of  carbon  monoxide  and  ferric  oxide,  manufacturers  increased  the 
dimensions  of  the  blast  furnaces,  particularly  in  England  and  America,  a 
height  of  thirty  meters  being  not  uncommon.  The  ratio  of  carbon  mon- 
oxide to  the  dioxide  in  the  escaping  gases  was  not  affected  however;  it 
was  thus  demonstrated  by  very  expensive  experience  that  the  reduction 
of  ferric  oxide  by  carbon  monoxide  has  a  limit.  A  study  of  the  laws  of 
chemical  equilibrium  would  have  led  to  this  conclusion  much  more 
quickly  and  above  all  much  less  expensively.  These  laws  teach  us  that: 

1 .  In  the  reduction  of  ferric  oxide  by  carbon  monoxide  an  equilibrium 
is  established  between  this  action  and  the  oxidation  of  iron  by  carbon 
dioxide. 

Fe2O3 + 3CO<=±2Fe  +  3C02. 


2.  The  ratio  CO:CO2  must  be  independent  of  the  pressure,  since  no 
change  in  the  volume  of  the  gas  takes  place  (§51). 

3.  This  ratio  varies  only  slightly  with  the  temperature,  since  very 
little  heat  is  generated  in  the  reaction. 

An  experimental  investigation  conducted  at  a  few  different  tempera- 
tures and  pressures  would  have  sufficed  to  determine  the  ratio  CO:CO2. 
The  result,  when  compared  with  the  ratio  CO:CO2  of  the  waste  gases, 
would  thus  have  shown  that  little  could  be  gained  by  an  increase  of  the 
furnace  dimensions.  This  illustrates  in  a  very  striking  way  the  value  of 
physical  chemistry  for  industrial  processes. 

Efforts  are  now  being  made  to  utilize  the  waste  gases  in  other  ways, 
such  as  by  burning  them  under  the  boilers  of  steam  engines  or  in  wind 
heaters  (for  heating  the  blast  air,  or  "wind").  In  recent  years  it  has 
been  found  that  greater  efficiency  is  attained  by  using  the  hot  waste- 
gases  directly  in  gas  engines  for  motive-power. 


466  INORGANIC    CHEMISTRY.  [§  304- 

304.  The  properties  of  iron  are  influenced  in  great  measure 
by  the  slight  admixtures  which  it  contains,  particularly  by  the 
carbon.  The  percentage  of  carbon  forms  the  ordinary  basis  of 
classification  of  the  different  grades  of  iron  under  the  heads>  pig 
iron  and  malleable  iron;  however,  in  the  industrial  world  this 
classification  is  not  always  adhered  to. 

Pig  iron,  or  cast  iron,  contains  2.3-5.1%  carbon.  It  fuses  very 
easily  but  there  is  no  previous  softening;  hence  it  is  not  malleable. 
It  is  brittle.  Pig  iron  is  the  direct  product  of  the  blast  furnaces 
and  the  iron  is  therefore  mixed  with  small  amounts  of  silicon,  phos- 
phorus, sulphur,  etc.  The  presence  of  manganese  makes  it  coarsely 
crystalline  and  it  is  then  known  as  spiegel-eisen.  This  is  utilized 
mainly  for  steel. 

Refined  iron,  containing  less  than  2.3%  carbon,  is  harder  to 
fuse,  but  is  extensible  and  malleable,  and  the  more  so  the  less  the 
impurities.  If  the  carbon  amounts  to  2.3-0.5%,  the  iron  can  be 
hardened;  in  this  manner  steel  is  obtained.  If  there  is  less  than 
0.5%  carbon,  it  can  no  longer  be  hardened;  this  is  wrought  iron. 
It  is  obvious  that  between  these  main  varieties  there  are  numerous 
intermediate  sorts,  which  are  prepared  in  such  a  way  as  to  suit  the 
purpose  for  which  they  are  intended. 

The  immense  commercial  importance  of  the  iron-carbon  system  has  led 
to  extensive  investigations  regarding  it,  notwithstanding  that  such  investiga- 
tions are  attended  by  great  experimental  difficulties,  partly  because  of  the 
very  high  temperatures  involved.  Because  of  these  difficulties  it  is  not  yet 
possible  to  give  an  entirely  satisfactory  representation  of  the  equilibrium 
conditions  concerned.  BAKHUIS  ROOZEBOOM,  CHARPY,  ROBERT-AUSTEN  and 
others  have  succeeded  in  working  out  the  accompanying  graphic  representa- 
tion which  indicates  the  behavior  of  the  system  in  the  main  at  least.  To 
appreciate  this  diagram  it  is  necessary  in  the  first  place  to  know  a  few  general 
facts  regarding  the  components  that  are  now  regarded  as  existing  in  the 
iron-carbon  system.  Distinction  is  made  between:  1.  ferrite,  or  chem- 
ically pure  iron  (pure  wrought  iron);  2.  martensite  (steel),  a  solid  solution 
of  carbon  in  iron.  It  is  so  regarded  because  microscopic  studies  have  shown 
that  martensite  is  always  homogeneous  in  spite  of  its  changing  carbon 
content,  which  may  be  as  high  as  2%.  3.  cementite  (the  commercial  white 
cast  iron),  an  iron-carbon  compound  of  the  formula  Fe3C;  and  4.  perlite> 
carboniferous  iron  (0.85%),  that  is  seen  under  a  high-power  microscope  to  be 
heterogeneous  and  is  regarded  as  a  eutectic  mixture  of  ferrite  and 
cementite. 

The  solidification  curve  of  a  binary  system  (§  237)  does  not  take  a  nor- 
mal course  in  the  iron-carbon  system.  Three  circumstances  complicate  the 


304.] 


IRON. 


467 


situation.  The  first  is  that  pure  iron  does  not  separate  out  of  the  molten 
mass,  but  that  we  obtain  the  solid  solution  martensite.  The  second  is  that 
changes  continue  to  occur  in  the  cooling  mass  after  complete  solidification; 
the  third  that  other  substances  separate  out  with  very  slow  cooling  than 
with  quick  cooling. 

We  may  consider  first  the  case  of  slow  cooling,  where  the  equilibria  that 
establish  themselves  between  solid  and  liquid  phases  are  presumably  stable: 
Let  us  assume  that  we  have  liquid  iron  with  a  carbon  content  below  4.3%. 
On  cooling  the  liquid  the  iron  begins  to  solidify  at  a  definite  temperature  (the 
point  G!  in  Fig.  74);  however,  it  is  not  pure  iron,  but  a  solid  solution  of 


1500 c 


(1000° 
G 


700' 


" 


FIG.  74. — IRON-CARBON  SYSTEM. 

carbon  in  iron  that  separates  out;  its  composition  is  shown  in  the  diagram 
by  the  point  dt.  If  the  carbon  content  of  the  fused  iron  is  a  different  one, 
we  have  separating  out  at  c2,  for  example,  the  solid  substance,  whose  com- 
position is  again  given  by  the  point  d2.  Thus  for  every  solidification  point 
of  the  curve  AC  we  can  find  a  point  dlf  d2,  etc.,  that  gives  the  composition 
of  the  solid  substance  which  begins  to  separate  out.  The  curve  AD  is 
the  geometrical  locus  of  these  points.  If,  therefore,  a  horizontal  line  is 
drawn  through  the  triangle  ADC,  the  point  c  gives  the  composition  of  the 
liquid  solution  which  solidifies  at  the  corresponding  definite  temperature 
(indicated  by  the  ordinate)  and  the  point  d  the  composition  of  the  solid 
solution  which  begins  to  separate  out  at  that  temperature.  At  C  the  eutectic 
point  is  reached.  Along  CB  graphite  separates  out;  at  C  itself  a  mixture 
of  graphite  and  martensite,  the  composition  being  given  by  D.  The  point 
C  is  at  1130°  and  4.3%  of  carbon.  The  martentite  formed  at  this  tem- 
perature contains  2%  of  carbon. 


468  INORGANIC   CHEMISTRY.  [§§304- 

Below  DC  all  is  solid;  but,  as  we  have  already  explained,  changes  con- 
tinue to  occur  in  the  solid  mass.  For  example,  if  martensite  is  heated,  it 
breaks  up  with  the  formation  of  graphite.  The  curve  DE  represents  the 
change  of  composition  of  the  solid  solution  with  falling  temperature  or,  in 
other  words,  it  represents  the  equilibrium  between  graphite  and  the  mixed 
crystals  (solid  solution)  at  different  temperatures.  Around  E,  where  the 
temperature  has  reached  about  700°,  the  martensite  contains  only  about 
0.85%  carbon.  At  the  point  E  the  formation  of  ferrite  begins. 

Finally,  the  curve  EG  indicates  the  composition  of  the  solid  solutions 
from  which  ferrite  separates  out.  Hence,  if  the  martensite  contains  less 
than  0.85%  of  carbon,  ferrite  is  deposited  along  EG,  exactly  as  ice  separates 
out  of  a  dilute  salt  solution  with  falling  temperature. 

If  the  cooling  is  sudden,  other  phases  are  formed,  the  limits  of  which 
are  represented  in  the  figure  by lines,  which  are  readily  under- 
stood. Instead  of  the  eutectic  point  C,  at  which  graphite  and  martensite 
separate  out,  we  have  a  eutectic  point  at  Cl}  very  close  to  C,  where  cementite, 
Fe3C,  separates  out  with  the  martensite.  Further,  the  line  C^E^  represents 
the  equilibrium  between  cementite  and  the  mixed  crystals  (martensite). 
At  El  ferrite  is  formed  together  with  cementite.  Martensite  changes  over 
at  this  temperature  into  a  eutectic  mixture  of  these  last  two  substances,  which 
has  the  fine  conglomerate  structure  so  characteristic  of  eutectics  and  is  known 

as  "  perlite."  Although  this  whole  system  shown  by lines  is  meta- 

stable,  it  can  exist  for  an  indefinite  period  at  ordinary  temperatures  because 
of  the  reduction  of  the  velocities  of  reactions  which  might  restore  the  stable 
forms. 

It  is  evident  from  the  above  that  with  slow  cooling  martensite  entirely 
disappears.  If  the  cooling  is  rapid,  however,  as  in  the  hardening  of  steel, 
martensite  can  be  brought  to  exist  at  ordinary  temperatures  even  though 
it  is  in  a  metastable  condition ;  its  transformation  velocity  is  then  extremely 
small.  If  the  hardened  steel  is  reheated,  it  changes  over  partially  into  the 
soft  conglomerate  of  ferrite  and  cementite;  this  is  what  takes  place  in  the 
"tempering"  of  steel. 

Small  admixtures  of  other  elements  have  an  effect  on  the 
properties  of  iron  equally  as  great  as  that  of  carbon.  The  presence 
of  silicon  has  about  the  same  effect  as  that  of  carbon,  but  it  is 
less  intense.  Sulphur  even  in  a  small  amount  renders  the  iron 
brittle  when  hot  and,  therefore,  useless  for  forging.  On  this 
account  sulphurous  ores  as  such  are  unsuitable  for  the  manu- 
facture of  iron.  Phosphorus  makes  the  iron  brittle  at  ordinary 
temperatures.  It  should  also  be  mentioned  that  as  a  general  rule 
the  effect  of  these  admixtures  is  strongly  modified  by  the  pres- 
ence of  others. 

305.  From  the  crude  pig  iron,  the  direct  product  of  the  blast 
furnace,  the  other  varieties  of  iron  are  prepared.  For  this  purpose 


305.] 


IRON. 


469 


it  must  be  freed  from  silicon,  sulphur,  phosphorus,  etc.,  as  well 
as  from  a  large  portion  of  its  carbon.  The  most  important  process 
for  accomplishing  this  commercially  is  the  BESSEMER  process.  The 
pig  iron  is  fused  and  run  into  a  pear-shaped  apparatus,  or  con- 
verter (Fig.  75),  in  the  bottom  of  which  are  holes  through  which 
air  is  blown  in.  Thus  by  the  oxidation  of  silicon,  manganese 
and  a  little  iron  and  without  the  use  of  fuel  the  temperature  is 
raised  high  enough  to  effect  the  burning  of  the  carbon.  The  BES- 
SEMER process  is  easier  controlled  if  the  elimination  of  carbon  is 
continued  past  the  steel  stage  and  until  molten  wrought  iron  is 
formed,  whereupon  enough  carboniferous  iron  is  added  to  furnish 


FIG.  75. — CONVERTER. 

steel  with  the  desired  percentage  of  carbon.     At  the  completion 
of  the  process  the  converter  is  emptied  by  tipping. 

In  some  European  mills  a  basic  converter  lining  containing  an  excess 
of  lime  and  magnesia  is  used.  The  phosphorus  in  the  ore  combines  with 
the  bases  to  form  phosphates,  which  enter  the  slag,  and  this  so-called 
"Thomas-slag"  is  used  in  large  quantities  as  a  fertilizer. 

The  only  successful  rival  of  the  BESSEMER  process  is  the 
SIEMENS,  or  open-hearth,  process.  By  employing  a  special  furnace 
and  gaseous  fuel  a  mixture  of  cast  iron  and  wrought  iron  (together 
with  some  iron  ore)  in  the  proper  proportions  can  be  fused  together 


470  INORGANIC  CHEMISTRY.  [§305. 

so  as  to  produce  a  very  good  steel.     A  basic  lining  can  also  be 
used  with  this  process. 

The  increased  demand  for  special  steels,  where  physical  and 
chemical  conditions  have  to  be  regulated  carefully,  has  given 
greater  significance  to  the  old  crucible  process,  the  steel  being 
made  in  graphite  crucibles  in  a  laboratory  manner  but  on  about 
ten  times  the  laboratory  scale.  Recently  electric  furnaces  of  the 
arc  and  induction  type  have  been  found  very  successful  in  pro- 
ducing (i  crucible "  steel  and  with  much  less  labor  than  the 


FIG.  76. — HEKOULT  FURNACE. 

crucible  process  requires.  The  cradle-shaped  HEKOULT  furnace 
is  shown  in  the  accompanying  combined  end-view  and  vertical 
section  (Fig.  76).  M  is  the  molten  metal,  S  the  slag,  and  E 
one  of  the  carbon  electrodes;  B  is  brick  lining  and  L  a  layer 
of  magnesium  silicate.  As  the  resistance  of  the  metal  is  small 
compared  with  that  of  the  slag  and  the  air,  most  of  the  heat 
is  generated  at  the  surface,  where  the  chemical  action  goes  on 
between  the  slag  and  the  metal.  The  furnace  is  eventually 
emptied  by  rocking  forward. 

Steel,   however  made,   is   a  very   complex   alloy,   containing 


305.]  IRON.  471 

carbon  and  manganese,  0.10-1.50%;  silicon  0.02-0.25%,  sulphur 
and  phosphorus  0.01-0.10%;  and  possibly  copper,  arsenic,  alu- 
minium, oxygen,  nitrogen,  and  cyanides,  and  is  capable,  as  has 
been  explained,  of  containing  the  iron  and  carbon  in  various 
combinations.  Steel  of  the  above  description  is  "  ordinary " 
steel.  Recently  a  large  market  has  developed  for  "  special " 
steels,  having  new  qualities,  especially  with  respect  to  hardness 
and  brittleness,  and  serving  new  purposes,  notably  in  tools, 
military  materials  and  materials  of  construction.  They  may  be 
produced  by  (1)  changing  the  physico-chemical  character  with 
respect  to  the  iron-carbon  system,  (2)  removing  harmful  occluded 
gases,  (3)  combining  other  elements  chemically  with  iron  or  carbon 
or  both,  and  (4)  adding  other  elements  to  form  isomorphous 
solutions  with  iron.  Steel  becomes  very  hard  and  brittle,  for 
instance,  when  it  is  suddenly  cooled  from  a  high  temperature. 
If,  however,  it  is  then  heated  for  a  definite  period  and  allowed 
to  cool  slowly,  it  becomes  more  or  less  tempered  according  to  the 
temperature,  i.e.,  it  can  be  made  to  have  any  desired  hardness 
and  elasticity  (within  certain  limits). 

Of  the  special  alloy  steels  the  nickel  steels,  chrome-nickel 
steels  and  chrome-vanadium  steels  seem  to  be  most  important. 
The  maximum  hardness  of  steel  is  reached  when  it  contains 
1-2%  carbon;  if,  however,  some  manganese  (up  to  8%)  or  chro- 
mium (up  to  1%)  is  added,  a  much  harder  modification  of  steel 
is  produced.  The  addition  of  nickel  gives  a  tougher  steel,  which 
is  especially  valuable  for  armor  plate.  Tungsten  (cf.  §  297)  and 
molybdenum  are  also  added  for  different  purposes.  In  any 
case,  however,  a  careful  heat  treatment  is  essential  to  develop 
the  desired  properties. 

The  production  of  wrought  iron  from  pig  iron  is  usually  accomplished 
by  the  puddling  process.  Pig  iron  is  melted  in  a  reverberatory  furnace 
lined  with  iron  ore  (oxide) ;  the  carbon  and  also  the  silicon  are  oxidized 
(and  so  removed)  partly  by  the  action  of  the  air,  but  mainly  by  that 
of  the  ore,  which  is  stirred  in  with  the  metal.  The  violent  reaction 
due  to  escaping  carbon  monoxide  gives  the  process  the  name  of  "pig- 
boiling."  The  iron  is  then  allowed  to  become  pasty,  when  it  is 
worked  up  into  large  masses  (blooms'),  which  are  removed  and  ham- 
mered and  rolled.  The  cinder  is  thus  squeezed  out  and  the  iron  is 
formed  into  bars. 


472  IXORGANIC  CHEMISTRY.  [§§  3Q5- 

Chemically  pure  iron  is  obtained  electrolytically  and  by  reduc- 
ing the  oxide  or  chloride  in  a  current  of  hydrogen.  If  the  reduc- 
tion takes  place  at  a  low  temperature,  the  resulting  iron  powder  is 
pyrophoric  (§  203).  It  is  a  silvery-white,  lustrous  metal  with  a 
specific  gravity  of  7.84  and  a  melting-point  as  high  as  1520°.  It 
boils  at  2450°.  It  is  the  most  magnetic  of  the  metals;  pure  iron 
and  wrought  iron  can  be  magnetized  only  temporarily;  steel, 
however,  permanently.  Iron  is  permanent  in  dry  air  or  in  water 
free  from  air  (CO2).  In  moist  air  it  rusts  rapidly  (§  279),  forming 
ferric  hydroxide;  as  the  rust  does  not  form  a  compact  film,  it 
keeps  on  forming. 

The  rusting  of  iron  is  greatly  retarded  by  contact  with  water  contain- 
ing a  little  alkali  or  salts  of  alkaline  reaction.  In  a  soda  solution,  for 
instance,  iron  remains  bright.  The  rusting  of  iron  in  contact  with 
water  can  be  explained  by  assuming  that  the  oxygen  dissolved  in  water 
endeavors  to  form  hydroxyl  ions  with  the  hydrogen  ions.  In  order  to 
compensate  their  negative  potential  the  iron  sends  its  positive  ions  into 
the  solution;  in  a  short  time  the  solubility  product  of  ferric  hydroxide 
is  reached  and  the  latter  is.  deposited;  in  other  words,  the  iron  rusts. 

Now,  if  hydroxyl  ions  are  previously  introduced  into  the  liquid  by 
the  addition  of  a  base  or  a  salt  of  alkaline  reaction,  the  ionization  of  the 
water  is  diminished  so  much  that  the  oxygen  can  find  almost  no  hydro- 
gen ions  with  which  to  form  hydroxyl  ions;  therefore  the  iron  does  not 
send  any  more  ions  (§§276  and  277)  into  the  solution  and  rusting  is 
greatly  retarded. 

Iron  dissolves  readily  in  hydrochloric  and  sulphuric  acids  with 
the  evolution  of  hydrogen.  At  red-heat  it  decomposes  water,  but 
the  oxide  is  also  reduced  by  hydrogen,  so  that  an  equilibrium 
results : 

3Fe  +  4H20<=»Fe3O4 + 4H2. 

In  nitric  acid  (not  too  concentrated)  iron  dissolves  readily  with 
the  evolution  of  nitric  oxide,  NO,  but  if  the  iron  is  first  dipped 
in  concentrated  nitric  acid  and  then  rinsed  off  it  becomes  indif- 
ferent to  the  action  of  nitric  acid.  This  so-called  "passivity" 
of  iron  is  probably  caused  by  a  very  thin  coating  of  oxide. 

Iron  forms  two  sets  of  salts,  the  ferrous  and  the  ferric. 


306.]  FERROUS  COMPOUNDS.  473 

Ferrous  Compounds. 

306.  In  the  ferrous  condition  iron  has  only  basic  properties. 

Ferrous  oxide,  FeO,  is  obtained  by  reducing  ferric  oxide  with 
carbon  monoxide.  It  is  a  black  powder,  which  oxidizes  easily 
on  warming.  Ferrous  hydroxide,  Fe(OH)2,  is  precipitated  from 
ferrous  salt  solutions  as  a  pale  green  gelatinous  substance  by  the 
addition  of  an  alkali;  it  oxidizes  very  rapidly  in  the  air  to  ferric 
hydroxide. 

Ferrous  chloride,  FeCl2,  is  formed  on  dissolving  iron  in  hydro- 
chloric acid;  it  crystallizes  from  this  solution  in  green  monoclinic 
prisms  containing  four  molecules  of  water.  The  anhydrous  salt 
is  obtained  as  a  white  sublimate  when  iron  is  heated  in  dry  hydro- 
chloric acid  gas.  With  potassium  chloride  and  ammonium  chlo- 
ride ferrous  chloride  forms  well  crystallized  double  salts,  e.g. 
FeCl2-2KCl  +  2H2O. 

Ferrous  sulphate,  FeS04+7H20  (green  vitriol,  copperas),  is  the 
most  familiar  ferrous  salt.  It  is  prepared  commercially,  princi- 
pally by  dissolving  up  the  waste  metal  of  steel- wire  factories  in 
sulphuric  acid,  but  also  by  partially  roasting  pyrite,  whereby  fer- 
rous sulphide,  FeS,  is  formed;  the  latter  is  left  exposed  to  the  ah-, 
when  it  oxidizes  gradually  to  ferrous  sulphate,  which  can  be  dis- 
solved out.  It  crystallizes  in  large,  bright  green,  monoclinic  prisms, 
which  effloresce  slightly  and  at  the  same  time  become  coated  with 
a  brown  layer  of  basic  ferric  sulphate.  The  double  salts  such  as 
FeS04-(NH4)2S04+6H2O,  MOHR'S  salt,  are  not  so  liable  to  oxi- 
dation; for  this  reason  use  is  frequently  made  of  MOHR'S  salt 
to  standardize  permanganate  solutions  (§  301).  Iron  vitriol  has 
numerous  uses,  e.g.  for  making  ink,  in  dyeing,  as  a  disinfectant 
(it  absorbs  both  ammonia  and  sulphuretted  hydrogen  and  is  there- 
fore used  to  dispel  bad  odors),  etc.,  etc. 

Ferrous  carbonate  is  somewhat  soluble  in  water  containing 
carbonic  acid  and  is  therefore  often  present  in  natural  waters 
(§  17).  The  basic  carbonate  which  is  precipitated  from  a  ferrous 
solution  by  soda  oxidizes  rapidly  in  the  air  to  ferric  hydroxide. 
The  latter  is  also  deposited  from  chalybeate  waters  on  standing  in 
the  air  for  a  time.  Ferrous  carbonate  is  only  known  as  a  mineral 
(siderite,  §  302). 


474  INORGANIC  CHEMISTRY. 


[§§307- 


Ferric  Compounds. 

3°7«  The  ferric  ion  has  only  very  slightly  basic  properties. 
Ferric  salts  of  weak  acids;  such  as  carbonic  acid,  do  not  exist.  In 
aqueous  solution  most  of  the  ferric  salts,  even  those  of  strong 
acids,  are  partially  hydrolyzed.  For  that  reason  they  are  brown- 
ish-red, since  this  is  the  color  of  ferric  hydroxide  in  colloidal  solu- 
tion. On  the  addition  of  an  excess  of  sulphuric  or  nitric  acid  this 
color  disappears,  because  there  is  no  longer  any  hydrolysis.  From 
this  it  appears  that  the  ferric  ion  itself  in  aqueous  solution  is  only 
slightly  colored.  The  ferric  salts  are  readily  converted  into  ferrous 
salts  by  reducing-agents. 

Ferric  oxide,  Fe20s,  iron  sesquioxide,  is  formed  on  heating 
various  iron  compounds  in  the  air  and  is  manufactured  by  igniting 
green  vitriol  (§  79).  It  is  a  dark-red  powder  and  finds  use  as  a 
pigment  (colcothar)  and  in  polishing  glass,  etc. 

Ferric  hydroxide  separates  out  as  a  reddish-brown  hydrogel, 
Fe203+ftH2O,  when  a  ferric  salt  solution  is  treated  with  an 
alkali.  The  freshly  precipitated  hydrogel  dissolves  in  a  solution 
of  ferric  chloride  or  acetate.  If  this  solution  is  dialyzed,  a  pure 
colloidal  solution  of  the  hydroxide  is  finally  obtained;  from  this 
the  hydrogel  is  reprecipitated  by  a  small  amount  of  alkali  or 
acid. 

Ferrous  ferric  oxide,  Fe3O4,  also  called  ferroso-ferric  oxide  or 
magnetic  iron  oxide,  occurs  in  nature  as  magnetite.  It  is  produced 
by  heating  iron  in  steam  (§  305). 

Ferric  chloride  is  obtained  by  passing  chlorine  into  a  solu- 
tion of  ferrous  chloride.  It  crystallizes  at  different  temperatures 
with  different  amounts  of  water,  being  an  example  of  the  case 
described  on  p.  341.  On  heating  the  salt  hydrochloric  acid  escapes 
with  the  water  of  crystallization.  Anhydrous  ferric  chloride  can 
be  prepared  by  heating  iron  in  a  current  of  dry  chlorine. 

Between  320°  and  440°  the  vapor  density  is  approximately  that 
calculated  for  Fe2Cl6;  between  750°  and  1050°  it  falls  to  half, 
indicating  a  splitting  off  of  chlorine  or  a  dissociation  into  2FeCl3. 

The  reddish-brown  color  of  the  aqueous  solution  of  ferric  chloride 
must  be  ascribed  chiefly  to  un-ionized  FeCls  molecules,  for  the  salt 
has  this  same  color  when  dissolved  in  ether,  in  which  no  ioniza- 
tion  occurs.  In  part,  also,  this  color  comes  from  ferric  hydroxide. 


308.]  FERRIC    COMPOUNDS.  475 

which  is  formed  by  hydrolytic  dissociation.  This  dissociation 
increases  on  warming  the  dilute  aqueous  solution,  for  a  very  dilute, 
almost  colorless  solution  of  ferric  chloride  turns  reddish-brown  on 
boiling.  When  cooled  the  liquid  gradually  resumes  its  original 
color. 

Ferric  sulphate,  obtained  by  dissolving  ferric  oxide  in  sulphuric 
acid,  forms  alums,  e.g.  potassium  iron  alum,  K2SC>4  ^62(864)  3+ 
24H20. 

When  a  ferrous  salt  is  converted  into  a  ferric  salt  in  aqueous 
solution  the  bivalent  ferrous  ion  is  transformed  into  a  trivalent 
ferric  ion.  The  oxygen  required  for  the  conversion  serves  to 
oxidize  the  hydrogen  ions  of  the  acid  (which  must  be  added)  to 
water,  whereupon  these  hydrogen  ions  surrender  their  charge  to 
the  iron  ions: 


10  +2(H'  +  C10  +  O  =2(Fe-  +  3Cl/)  +H2O. 

Ferrous  chloride         Hydrochloric  acid  Ferric  chloride 

Inversely,  the  reduction  of  ferric  salts  to  ferrous  salts  can  be 
explained  by  supposing  that  every  ferric  ion  gives  up  a  third  of  its 
charge  to  another  atom  and  thus  makes  the  latter  an  ion  or  neu- 
tralizes its  charge. 

Salts  of  iron  are  also  known  which  are  derived  from  the  hypothetical 
oxide  FeO3.  They  are  obtained  by  heating  iron  filings  with  saltpetre  or 
passing  chlorine  into  an  alkaline  suspension  of  the  ferric  oxide  hydrogel. 
From  such  solutions  potassium  ferrate,  K2FeO4,  crystallizes  out  in  dark- 
red  prisms,  isomorphous  with  the  chromate  and  sulphate  of  potassium. 
These  crystals  are  readily  soluble  in  water,  but  their  dark-red  solution 
soon  decomposes  with  the  separation  of  ferric  hydroxide  and  oxygen 
gas.  The  free  ferric  acid  is  unknown. 

308.  Iron  unites  with  cyanogen  to  form  complex  and 
unusually  stable  anions,  viz.,  the  jerrocyanic  ion  [Fe(CN)6]""  and 
the  ferricyanic  ion  [Fe(CN)6]'".  Their  best-known  salts  are  potas- 
sium ferrocyanide,  K4Fe(CN)6-3H2O,  and  potassium  ferricyanide, 
K3Fe(CN)6,  the  yellow  and  red  prussiates  of  potash, 
respectively.  The  ionization  of  the  complex  ions  themselves  is  so 
slight  that  they  give  none  of  the  ordinary  reactions  for  iron. 

For  the  commercial  manufacture  of  yellow  prussiate  of  potash 
two  processes  are  used:  .In  the  first,  animal  refuse  (e.g.  blood) 
is  charred,  yielding  a  black,  highly  nitrogenous  mass.  This  is 


476  INORGANIC   CHEMISTRY.  [§§308- 

ignited  with  potash  and  iron  filings.  After  cooling,  hot  water  is 
added  and  the  mixture  filtered;  from  this  filtrate  the  yellow 
prussiate  crystallizes  out  on  standing.  This  salt  is  not  formed 
until  the  ignited  mass  is  treated  with  water,  for  yellow  prussiate  is 
decomposed  by  heat  and  cannot  therefore  be  present  in  the  ignited 
mass.  The  latter  probably  contains  potassium  cyanide,  iron  and 
iron  sulphide  (animal  refuse  always  contains  sulphur  compounds')  . 
These  substances  can  interact  according  to  the  equations: 

6KCN+FeS  =K4Fe(CN)6+K2S; 


Fe(CN)2  +  4KCN  =  IQFe(CN)«. 

The  second  process  is  employed  in  illuminating-gas  factories, 
for  the  unpurified  gas  contains  a  little  cyanogen  and  prussic  acid. 
After  being  freed  from  tar  and  ammonia  it  is  passed  through  a 
washer  (scrubber)  containing  a  solution  of  potash  in  which  ferrous 
carbonate  (ferrous  sulphate  +  potassium  carbonate)  is  suspended. 
The  following  reactions,  among  others,  are  known  to  go  on  here: 

FeCO3  +  2HCN<=±Fe(CN)2  +  H2O  +  CO2; 

K2CO3  +  2HCN^2KCN  +  H2O  +  CO2. 

Notwithstanding  that  these  reactions  are  reversible,  the  hydro- 
cyanic acid  can  be  quantitatively  fixed  in  this  way,  because  the 
ferrous  cyanide  and  potassium  cyanide  interact  to  form  potassium 
ferrocyanide,  which  is  but  very  slightly  affected  by  carbon  dioxide. 

Potassium  ferrocyanide,  K4Fe(CN)6-3H2O,  forms  large  sulphur- 
colored  crystals.  Its  three  molecules  of  water  can  be  expelled  by 
gently  warming,  whereupon  the  salt  is  left  as  a  white  powder.  It 
is  not  poisonous.  With  dilute  sulphuric  acid  it  produces  prussic 
acid  on  warming;  with  concentrated  sulphuric  acid  it  yields 
carbon  monoxide. 

The  free  ferrocyanic  acid,  H4Fe(CN)6,  separates  out  as  a  white 
crystalline  precipitate  when  concentrated  hydrochloric  acid  is 
added  to  a  strong  solution  of  potassium  ferrocyanide.  The  pre- 
cipitate soon  turns  blue  in  the  air  on  account  of  the  formation  of 
Prussian  blue  (and  partial  decomposition  as  well).  Various  salts 
of  this  acid  have  characteristic  colors  and  are  insoluble;  hence 
potassium  ferrocyanide  finds  use  in  analysis.  It  is  an  interesting 


309.]  COBALT  AND  NICKEL.  477 

fact  that  this  compound  of  iron  can  serve  as  a  distinguishing 
reagent  for  ferrous  and  ferric  compounds.  The  ferrous  salt  of 
ferrocyanic  acid  is  white,  but  in  the  presence  of  air  it  passes  rapidly 
over  into  the  blue  ferric  salt  (Prussian  blue—a  valuable  pigment). 
The  copper  salt  (§  40)  is  brownish-red,  the  zinc  salt  white,  etc. 

Sodium  nitroprusside,  Na2Fe(CN)5(NO)  -2H2O,  is  formed  by 
the  action  of  nitric  acid  on  sodium  ferrocyanide.  It  crystallizes 
in  ruby-red  prisms  and  is  a  delicate  reagent  for  alkali  sulphides, 
whose  solutions  it  colors  violet. 

Potassium  ferricyanide,  K3Fe(CN)6,  red  prussiate  of  potash,  is 
formed  from  the  yellow  prussiate  by  treating  a  solution  of  the 
latter  with  chlorine  or  bromine : 

K4Fe(CN)6  +C1  =  KC1 +K3Fe(CN)6. 

It  appears  in  dark-red  crystals,  which  are  readily  soluble  in  water. 
The  aqueous  solution  is  unstable.  The  salt  is  often  employed  as  an 
oxidizing-agent  in  alkaline  solution,  being  itself  converted  into 
the  ferrocyanide: 

2K3Fe(CN)6  +2KOH  ==2K4Fe(CN)6  +H20  +0. 

Iron  forms  some  very  peculiar  compounds  with  carbon  monoxide: 
Fe(CO)4  and  Fe(CO)5.  They  are  produced  when  carbon  monoxide  is 
passed  over  finely  divided  iron  at  80°,  or  at  ordinary  temperatures  if 
the  gas  is  under  pressure.  Iron  vessels  which  have  held  compressed 
illuminating-gas  for  some  time  are  more  or  less  attacked  by  the  carbon 
monoxide  of  the  gas,  for  if  gas  which  has  been  kept  in  such  a  vessel  is 
allowed  to  escape  through  a  hot  glass  tube  an  iron  mirror  is  formed  on 
the  inside  of  the  tube. 

COBALT  AND  NICKEL. 
Cobalt. 

309.  The  two  best-known  minerals  of  this  metal  are  smaltite, 
CoAs2,  and  cobaltite,  or  cobalt  glance,  CoAsS.  The  metal  is  ob- 
tained by  calcining  these  minerals  and  reducing  the  resulting 
cobalto-cobaltic  oxide,  Co304,  with  carbon  (or  hydrogen).  It  has 
a  pink  color  and  a  high  lustre.  Sp.  g.  8.9;.  m.-pt.,  1490°.  It 
is  magnetic  but  much  less  so  than  iron.  It  is  indifferent  to  the 
air.  Hydrochloric  and  sulphuric  acids  dissolve  it  very  slowly 
but  it  readily  forms  a  nitrate  with  nitric  acid. 


478  INORGANIC   CHEMISTRY.  [g§  300- 

Besides  the  oxide,  Co304,  just  referred  to  there  are  two  others, 
cobaltous  oxide,  CoO,  and  cobaltic  oxide,  Co2O3.  The  salts  are  all 
cobaltous,  corresponding  to  the  bivalent  ion  Co". 

COBALTOUS  COMPOUNDS. 

The  solutions  of  the  salts  are  red;  hence  this  is  the  color  of 
the  cobalt  ion.  The  non-ionized  cobalt  salts  are  blue,  e.g.,  the 
anhydrous  CoCl2,  the  silicate,  etc.  This  difference  in  color  enables 
us  to  tell  readily  whether  a  cobalt  salt  in  solution  is  ionized  or 
not.  Thus  in  concentrated  solutions,  for  instance,  all  those  cir- 
cumstances which  reduce  the  ionization  cause  a  change  of  color 
from  red  to  blue,  e.g.  when  a  concentrated  cobalt  chloride  solu- 
tion is  warmed  or  treated  with  hydrochloric  acid.  That  the  ioni- 
zation is  diminished  by  warming  was  mentioned  in  connection  with 
cupric  chloride  (§  244). 

Cobaltous  chloride,  CoCl2-6H20,  forms  red  monoclinic  crys- 
tals, which  turn  blue  on  heating  because  of  dehydration.  Cobalt 
sulphate,  CoSCU  -7H20,  is  obtained  in  dark-red  monoclinic  prisms 
and  is  isomorphous  with  FeS04-7H20.  It  forms  double  salts 
with  alkali  sulphates,  e.g.  K2S04-CoS04  +  6H20.  Cobalt  nitrate, 
Co(N03)2-6H20,  appears  in  red  hygroscopic  prisms.  Cobalt  sili- 
cate is  very  deep  blue;  hence  its  use  for  coloring  glass.  Pulverized 
cobalt  silicate  serves  as  a  pigment  (smalt)  in  painting,  etc.  TH£- 
NARD'S  blue  is  a  pigment,  obtained  by  igniting  cobalt  salts  with 
alumina. 

COBALTIC  COMPOUNDS. 

310.  Cobaltic  oxide,  Co203,  is  obtained  by  igniting  cobalt 
nitrate.  It  is  a  black  powder,  which  passes  over  into  cobalto- 
cobaltic  oxide,  00364,  at  red  heat  and  at  white  heat  yields  cobaltous 
oxide.  It  has  the  character  of  a  peroxide;  for  by  the  addition  of 
sulphuric  acid  it  is  converted  into  a  cobaltous  salt  with  the  evolu- 
tion of  oxygen  and  it  yields  chlorine  with  hydrochloric  acid.  How- 
ever, in  cold  dilute  hydrochloric  acid  it  dissolves  without  generat- 
ing scarcely  any  chlorine. 

Like  iroa,  cobalt  also  forms  complex  ions,  of  which  those  with 
cyanogen  are  very  stable.  There  are  cobalt  salts  corresponding  in 
composition  to  the  yellow  and  the  red  prussiates  of  potash;  the 
salt  Iv3Co(CN)6,  potassium  cobalticyanide,  crystallizes  in  colorless 


3ii.[  cu/tM.r  A.VD  \ICKI:L.  479 

rhombic  prisms.  A  peculiar  complex  ion  occurs  in  the  potassium 
cobaltic  nitrite,  6KNO2-Co2(N02)6+nH20,  or  K3.Co(NO2)6  + 
wH20.  It  is  formed  on  treating  a  solution  of  a  cobalt  salt,  with 
potassium  nitrite  and  acetic  acid.  It  is  a  yellow  crystalline  pre- 
cipitate, which  is  very  slightly  soluble  when  potassium  ions  are 
present,  in  excess  in  the  liquid. 

Cobalt  also  forms  numerous  complex  ions  with  ammonia  (§317). 

Nickel. 

311.  Nickel  occurs  in  niccolitc,  NiAs,  and  nickel  glance,  or 
f/rr.sWo/7///r,  XiAsS.  Ks|)cci;illy  iiu| >nrl ;i nl  is  the  nickel  silic;ite, 
garnierite,  H2(Ni,Mg)Si04-faq(?),  which  was  discovered  by  GAB- 
NIER  in  New  Caledonia,  where  it  occurs  in  enormous  quantities. 
Canada,  too,  has  some  rich  nickel  deposits.  From  this  ore  the 
nickel  is  obtained  by  a  blast-furnace  process  similar  to  that  for 
iron.  The  discovery  of  garnierite  marked  the  beginning  of  a 
new  era  in  the  nickel  industry.  Much  nickel  is  refined  electro, 
lytically. 

Nickel  is  almost  as  white  as  silver,  is  very  tough  and  has  a 
high  metallic  lustre.  Sp.  g..=  8.8-9.1;  m.-pt.  - 1452°.  It  is 
feebly  magnetic.  It  dissolves  sparingly  in  hydrochloric  and 
sulphuric  acids  but  freely  in  nitric  acid.  It  is  permanent  in  the 
air. 

It  is  employed  in  nickel-plating  metallic  objects  and  as  a  con- 
stituent of  several  alloys.  German  silver  contains  about 
50%  copper,  25%  nickel,  and  25%  zinc.  The  nickel  coins  of 
Germany  and  the  United  States  consist  of  75%  copper  and  25% 
nickel.  The  use  of  nickel  to  vary  the  properties  of  iron  has  already 
been  mentioned  (§  305). 

The  oxides  of  nickel,  NiO  and  Ni20a,  are  very  similar  to  those 
of  cobalt.  The  nickelous  oxide,  NiO,  is  the  only  one  which  forms 
s;i,lts. 

Nickel  chloride,  NiCl2-6H20,  yields  green  monoclinic  prisms. 
When  healed  it  turns  yellow  on  account  of  loss  of  water. 

Nickel  sulphate,  NiSO4-7H20,  crystallizing  in  green  rhombic 
prisms,  is  isomorphous  with  the  corresponding  ferrous  and  other 
salts  and  also  forms  analogous  double  salts. 

Nickelic  oxide,  Ni203,  also  behaves  as  a  peroxide;  when  warmed 
with  hydrochloric  acid  it  yields  chlorine  gas  and  nickel  chloride. 


480  INORGANIC  CHEMISTRY.  [§§311- 

Nickel  carbonyl,  Ni(CO)4,  is  formed  when  carbon  monoxide  is 
led  over  finely  divided  nickel  at  ordinary  temperatures.  A  state 
of  equilibrium  results  here,  viz.: 

Ni+4CO<=±Ni(CO)4, 

which  is  displaced  to  the  left  with  rising  temperature,  since  the 
decomposition  of  nickel  carbonyl  takes  place  with  a  considerably 
absorption  of  heat  (§  103).  Even  as  low  as  60°  the  decomposition 
is  of  an  explosive  nature.  It  follows  from  the  above  equation  that 
an  increase  of  pressure  (§  122)  must  greatly  increase  the  propor- 
tion of  nickel  carbonyl  formed.  Experiments  confirming  this 
showed  at  the  same  time  that  both  the  formation  and  the  decom- 
position of  this  compound  are  very  sensitive  to  traces  of  foreign 
substances. 

Nickel  carbonyl  is  a  colorless,  highly  refractive  liquid,  which 
boils  at  43°  and  congeals  (crystalline)  at  —23°.  When  heated  in 
the  air  it  burns  with  a  very  sooty  flame.  This  compound  is  of 
aid  in  extracting  nickel  from  low-grade  ores. 

Nickel  also  forms  a  complex  ion  with  cyanogen.  On  dissolving 
nickel  cyanide  in  an  excess  of  potassium  cyanide  the  compound 
K2Ni(CN)4is  produced;  it  is,  however,  unstable,  being  decomposed 
by  hydrochloric  acid  with  the  deposition  of  nickel  cyanide,  Ni(CN)2. 

312.  A  peculiar  property  is  exhibited  by  the  sulphides  of  cobalt 
and  nickel,  CoS  and  NiS.  Hydrogen  sulphide  does  not  precipitate 
these  sulphides  from  acid  solutions,  but,  once  precipitated  (by 
ammonium  sulphide),  they  are  apparently  not  redissolved  by  dilute 
acids.  This  is  contrary  to  the  general  rule  of  §  146  (see  also  §  73), 
for  the  sulphide  should  either  be  precipitated  by  hydrogen  sulphide 
from  a  feebly  acid  solution  (e.g.  CuS),  which  is  the  case  when  the 
solubility  product  is  very  small,  or  else,  when  the  solubility  product 
is  larger,  it  should  dissolve  in  dilute  acids,  as  is  the  case  with  ferrous 
sulphide.  As  a  matter  of  fact,  however,  no  real  anomaly  exists 
here,  for  the  rate  of  solubility  of  these  sulphides  is  only  very  slow 
under  the  usual  conditions  of  the  reaction,  viz.,  dilute  acid  and 
room  temperature.  It  increases  with  the  concentration  of  the  acid, 
temperature  of  reaction  and  fineness  of  grain  of  the  precipitate. 
Nickel  sulphide  is  soluble  in  alkali  sulphides  immediately  upon  its 
formation,  but  when  once  deposited  in  the  solid  state  it  is  insol- 
uble, or  nearly  so.  This  is  seen  when  a  nickel  solution  is  treated' 


313.] 


PLATINUM    METALS. 


431 


with  tartaric  acid  and  then  with  an  excess  of  sodium  hydroxide, 
no  nickelous  hydroxide  being  precipitated.  If  hydrogen  sulphide 
is  passed  into  this  solution  a  very  dark-colored  liquid  results,  from 
which  nickel  sulphide  is  deposited  only  very  slowly.  The  same  is 
true  of  cobalt  in  very  dilute  solution;  in  concentrated  solution, 
however,  the  cobalt  sulphide  soon  passes  over  into  the  insoluble 
modification  and  separates  out. 


PLATINUM  METALS. 

313.  Under  this  head  are  included  the  metals  ruthenium,  rho- 
dium, palladium,  osmium,  iridium,  and  platinum.  They  occur 
only  in  metallic  form  and  are  associated  together  in  mixtures 
or  combinations.  The  principal  deposits  are  in  the  Ural  and 
Caucasus,  but  smaller  quantities  are  also  found  in  Colombia, 
Brazil  and  Borneo.  The  Ural  yields  95%  of  the  total  production. 
The  most  important  of  these  metals  is  platinum.  The  platinum 
ores  usually  contain  admixtures  of  iron,  gold,  etc. 

This  group  falls  into  two  subdivisions :  the  light  metals, 
ruthenium,  rhodium  and  palladium,  and  the  heavy  metals, 
osmium,  iridium  and  platinum.  The  two  sub-groups  differ  con- 
siderably in  atomic  weight  and  specific  gravity: 


Light. 

Heavy. 

Ru 

Rh 

Pd 

Os 

Ir 

Ft 

Atomic  weight.  .  .  . 
Specific  gravity.  .  . 

101.7 
12.26 

103.0 
12.1 

106.5 
11.9 

191 
22.4 

193.0 
22.38 

194.8 
21.45 

A  complete  separation  of  the  platinum  metals  from  each  other 
is  extremely  difficult,  in  the  first  place  because  their  properties  are 
very  similar,  and  in  the  second  place  because  their  behavior  is  con- 
siderably modified  by  their  mutual  presence — a  fact  which  indi- 
cates the  existence  of  compounds  with  each  other.  Thus,  for 
instance,  platinum  dissolves  readily  in  aqua  regia  while  pure  iridium 
is  insoluble  in  it;  nevertheless  when  an  alloy  of  the  two  metals  is 
treated, with  aqua  regia,  some  of  the  iridium  is  carried  into  solu- 


482  INORGANIC  CHEMISTRY.  [§§313- 

tion.  Further,  the  presence  of  iron  (which  occurs  in  all  platinum 
ores)  is  often  very  disturbing;  for  example,  pure  platinum  solu- 
tions are  not  precipitated  by  soda  or  barium  carbonate,  but  if  iron 
is  present  more  or  less  platinum  comes  down  with  the  hydroxide  of 
iron.  In  spite  of  these  difficulties  platinum,  palladium,  rhodium 
and  iridium  can  now  be  purchased  in  a  remarkably  pure  state. 

For  the  manufacture  of  platinum  and  the  other  metals  of  the  group 
in  the  pure  state  the  various  factories  employ  their  own  secret  methods. 
In  general  the  procedure  is  about  as  follows:  The  ore  is  first  treated 
with  aqua  regia  to  dissolve  out  the  major  part  of  the  "noble"  metals, 
leaving  in  the  residue  the  alloy  iridosmine,  besides  more  or  less  sand. 
The  ore  thus  consists  not  of  one  alloy,  but  of  two:  the  crude  platinum 
and  the  iridosmine.  Both  contain  all  six  platinum  metals,  although  in 
different  relative  amounts.  The  crude  platinum  contains,  besides 
platinum,  principally  palladium,  rhodium,  and  iridium,  while  the  iridos- 
mine, as  its  name  indicates,  consists  mainly  of  iridium  and  osmium. 

It  is  comparatively  easy  to  separate  out  the  osmium  and  ruthenium; 
they  form  volatile  oxygen  compounds  and  can  therefore  be  removed 
by  distillation.  Platinum  and  iridium  give  difficultly  soluble  com- 
pounds with  ammonium  chloride,  which  in  turn  are  reduced  to  the 
metal  form  by  ignition.  However,  if  the  solution  of  the  crude  platinum 
is  precipitated  with  ammonium  chloride,  the  resulting  precipitate  is 
found  to  contain  considerable  amounts  of  rhodium  and  palladium,  and 
an  extended  procedure  is  necessary  for  the  isolation  of  the  pure  metals. 
On  the  other  hand,  it  is  not  possible  to  precipitate  the  platinum  and 
iridium  completely  in  this  way ;  the  filtrate  from  the  ammonium  chloride 
contains  palladium  and  rhodium,  with  smaller  amounts  of  iridium  and 
platinum,  and,  in  order  to  work  it  up,  a  further. complicated  procedure 
is  necessary. 

Ruthenium. 

314.  This  steel-gray  metal  occurs  only  in  very  small  quantities;  it 
is  hard,  very  brittle,  and  very  difficult  to  fuse,  a  temperature  of  at  least 
1800°  being  necessary.  Even  when  finely  divided  it  is  but  very  sparingly 
soluble  in  aqua  regia,  forming  Ru2Cl6,  but  when  alloyed  with  platinum, 
it  dissolves  readily.  The  compound  RuCl4  is  known  only  in  double  salts. 
As  a  powder  the  metal  oxidizes  in  the  air  to  RuO  and  Ru203.  Ruthe- 
nium also  forms  characteristic  salts,  in  which  it  plays  the  part  of  an 
acid. 

Potassium  ruthenate,  K2RuO4,  results  from  fusing  ruthenium  with 
caustic  potash  and  saltpetre.  It  crystallizes  with  !H2Oin  black  prisms  of 


314.]  PLATINUM  METALS.  483 

a  greenish  lustre.  With  water  it  forms  a  dark  orange-colored  solution. 
Its  conduct  reminds  one  of  potassium  manganate,  for  under  the  influence 
of  dilute  acids  it  is  converted  into  potassium  perruthenate,  KRu04,  with 
the  simultaneous  precipitation  of  a  black  oxide,  Rh2O5  (or  RuO2?).  It 
crystallizes  in  black  octahedrons  of  metallic  lustre,  which  dissolve  in 
water  to  a  dark-green  solution.  A  peculiar  compound  is  the  tetroxide, 
RuO4,  which  volatilizes  when  chlorine  is  passed  into  the  concentrated 
solution  of  potassium  ruthenate.  It  can  be  solidified  by  cooling,  when 
it  forms  a  golden  crystalline  mass,  fusible  at  25.5°.  There  is  no  acid 
corresponding  to  this  oxide.  RuO4  is  used  for  the  preparation  of  pure 
ruthenium. 

Osmium. 

This  metal  is  very  analogous  to  ruthenium;  it  melts  as  high  as  2500°. 
The  chlorides  OsCl2  and  OsCl4  and  the  oxides  OsO,  Os2O3  and  OsO2  are 
known.  The  great  similarity  to  ruthenium  is  especially  noticeable  in 
the  highest  oxides.  Thus  fusion  with  caustic  potash  and  saltpetre  pro- 
duces potassium  osmiate,  K2OsO4,  which  crystallizes  from  aqueous  solu- 
tion in  dark-violet  octahedrons  containing  two  molecules  of  water. 
The  characteristic  osmium  compound  is  the  tetroxide,  OsO4,  formed 
by  igniting  finely  powdered  osmium  in  the  air  or  by  the  action  of  chlo- 
rine on  the  metal  in  the  presence  of  water.  The  aqueous  solution  of 
Os04  reacts  neutral,  but  is  often  (wrongly)  called  osmic  acid.  It  is 
employed  in  microscopy  since  organic  substances  (i.e.  reducing-agents) 
reduce  it  to  black  osmium.  No  salts  derived  from  OsO4  are  known. 
This  compound  is  used  in  preparing  pure  osmium. 


Rhodium. 

The  metal  in  the  fused  state  has  the  appearance  of  aluminium  and 
is  just  as  extensible  (malleable  and  ductile)  as  silver.  It  is  prepared 
pure  in  the  arts  by  way  of  the  chloro-purpureo  rhodium  chloride, 
Rh(NH3)5Cl3  (cf.  §  317).  Neither  acids  nor  aqua  regia  affect  it.  When 
heated  in  the  air  it  is  oxidized  to  the  rhodious  oxide,  RhO.  It  is  able  to 
absorb  a  considerable  amount  of  hydrogen.  The  rhodic  oxide,  Rh2O3, 
yields  salts  with  acids.  Of  the  chlorides  only  Rh2Cle  is  known;  this  is 
obtained  by  direct  synthesis  as  a  reddish-brown  substance;  it  forms  sol- 
uble double  salts  with  the  alkali  chlorides.  The  most  satisfactory 
thermocouple  for  measuring  high  temperatures  is  ma'de  of  pure  platinum 
and  an  alloy  of  platinum  and  rhoplium. 


484  INORGANIC  CHEMISTRY.  [§§314- 

Iridium. 

This  very  refractory  metal  is  obtained  from  iridosmine  by  heating  in  a 
current  of  oxygen,  when  the  osmium  volatilizes  as  tetroxide.  In  the  form 
of  a  platinum  alloy  it  is  employed  in  the  manufacture  of  "platinum" 
crucibles,  dishes,  distilling  vessels  for  the  concentration  of  sulphuric  acid 
(§  86),  etc.  The  prototype  of  the  meter  at  Paris  is  made  of  an  alloy 
of  90%  platinum  and  10%  iridium.  The  admixture  of  iridium  makes  the 
platinum  more  indifferent  to  chemical  agents,  although  at  high  tem- 
peratures the  volatility  of  the  iridium  is  often  troublesome.  When  pure, 
iridium  is  not  attacked  by  aqua  regia. 

Iridium  forms  two  chlorides,  Ir2Cl6  and  IrCl4.  Both  of  them  give 
double  salts  with  the  alkali  chlorides;  e.g.  Ir2Cl6  •  6KC1  +  6H2O  and 
IrCl4-2KCl.  The  former  dissolves  in  water  readily,  the  latter  with 
difficulty.  The  tetrachloride  appears  as  a  black  substance  forming 
with  water  an  intensely  red  solution.  For  this  reason  a  platinum  chlo- 
ride solution  which  contains  iridium  has  a  much  deeper  color  than  a 
pure  solution. 

Palladium. 

315.  The  silvery-white  metal  fuses  at  1549°,  i.e.  more  easily  than 
platinum.  When  finely  divided  it  dissolves  in  boiling  concentrated 
hydrochloric,  sulphuric  and  nitric  acids.  On  ignition  in  the  air  it  is 
at  first  oxidized,  thus  losing  its  lustre,  but  at  a  higher  temperature  the 
metallic  lustre  reappears.  The  most  peculiar  characteristic  of  the 
metal  is  its  ability  to  absorb  hydrogen  in  large  quantities  (occlusion). 
Freshly  ignited  palladium  foil  absorbs  370  times  its  own  volume  of 
hydrogen  at  room  temperature.  By  making  palladium  foil  the  cathode 
in  a  water  electrolysis  apparatus  the  metal  can  be  made  to  take  up 
even  960  times  its  own  volume.  This  absorption  does  not  alter  its 
metallic  appearance.  The  absorbed  hydrogen  can  all  be  expelled  by 
heating  in  a  vacuum. 

Palladium  charged  with  hydrogen  is  a  strong  reducing-agent ;  chlo- 
rine and  iodine  are  reduced  by  it  (see  §  200)  to  hydrogen  chloride  and 
hydrogen  iodide,  respectively,  and  ferric  salts  are  reduced  to  ferrous  salts. 

Palladium  forms  two  series  of  compounds,  the  -ous  PdX2,  and  the 
-ic,  PdX4.  A  characteristic  compound  of  the  first  series  is  palladious 
iodide,  PdI2,  which  is  precipitated  by  potassium  iodide  from  solutions  of 
-ous  salts  as  a  black  insoluble  substance.  This  reaction  is  occasionally 
used  to  separate  iodine  from  the  other  halogens,  since  their  palladium 
compounds  are  readily  soluble. — Palladia  chloride,  PdCl4,  is  produced 
by  dissolving  the  metal  in  aqua  regia.  With  KC1  or  NH4C1  it  forms 
a  difficultly  soluble  double  chloride,  K2PdCl6  or  (NH4)2PdCl6.  On  the 
evaporation  of  its  solution  PdCl4  dissociates  into  PdCl2.  and  C12. 


316.]  PLATINUM   METALS.  485 

Platinum. 

316.  This  metal  is  the  principal  constituent  of  the  platinum 
ores.  It  fuses  at  about  1760°  and  is  extremely  malleable  and 
ductile,  hence  it  can  be  made  into  very  fine  wire  and  very  thin 
foil.  Since  it  becomes  soft  at  red  heat  it  can  be  easily  worked. 
Platinum  is  used  in  the  greatest  variety  of  ways,  among  others 
the  manufacture  of  utensils  for  the  chemical  laboratory,  in  the 
distillation  of  sulphuric  acid,  in  electrical  apparatus,  in  jewelry 
for  settings  of  gems,  in  electric  furnaces,  in  incandesecent  lights 
and  dentistry.  The  last  named  use  consumes  about  one-third 
of  the  total  production.  When  finely  divided  it  absorbs  oxygen 
(partially  combining  with  it),  a  property  to  which  is  attributed  the 
phenomenon  that  numerous  oxidations  proceed  with  unusual  ease 
in  the  presence  of  platinum.  Its  use  in  gas-lighters  depends  on  this 
property.  When  the  metal  is  precipitated  from  its  solutions  by 
reducing-agents,  it  is  frequently  obtained  as  an  extremely  fine 
velvet-black  powder,  platinum  black.  When  the  double  chloride 
(NH4)2PtCl6  is  ignited,  the  metal  is  left  as  a  porous  mass — platinum 
sponge.  At  red  heat  a  platinum  partition  allows  hydrogen  to  pass 
through,  while  other  gases  are  held  back.  This  is  due  to  the  forma- 
tion of  a  compound  or  to  the  solubility  cf  hydrogen  in  platinum. 
Various  substances  attack  platinum  at  elevated  temperatures,  e.g. 
the  hydroxides,  cyanides  and  sulphides  of  the  alkalies;  hence  these 
substances  should  not  be  fused  in  platinum  vessels.  This  also 
applies  to  lead  and  other  heavy  metals,  for  they  form  low-melting 
alloys  with  platinum. 

There  are  two  sets  of  platinum  compounds  according  to  the  gen- 
eral formula  PtX2  and  PtX4.  The  best-known  platinum  compound 
is  chlorplatinic  acid,  H2PtCl6,  obtained  by  dissolving  platinum  in 
aqua  regia.  When  the  solution  is  evaporated  the  chlorplatinic 
acid  is  left  in  the  form  of  large,  reddish-brown,  very  hygroscopic 
prisms.  Its  aqueous  solution  contains  the  anion  PtCle",  for  such 
an  anion  goes  to  the  anode  in  an  electrolysis;  silver  nitrate 
precipitates  from  the  solution  not  silver  chloride,  which  it  would 
certainly  do  if  free  chlorine  ions  were  present,  but  the  compound 
Ag2PtCls.  Two  characteristic  salts  of  this  acid  are  those  of  potas- 
sium and  ammonium;  they  are  very  difficultly  soluble  in  water 
and  insoluble  in  alcohol;  when  the  aqueous  solution  is  evaporated 


486  INORGANIC   CHEMISTRY.  [§§316- 

the  salt  remains  in  the  form  of  small,  but  well-formed  octa- 
hedrons of  a  golden  hue.  The  potassium  salt  is  often  made 
use  of  in  determining  potassium  when  sodium  is  also  present, 
the  sodium  platinic  chloride  being  very  soluble,  even  in  alcohol. 

Of  the  remaining  platinum  compounds  a  few  may  be  referred 
to.  If  a  solution  of  the  above  acid,  H2PtCl6,  is  treated  with 
sodium  hydroxide  and  then  with  acetic  acid,  platinum  hydroxide, 
Pt(OH)4,  is  precipitated.  It  is  soluble  in  strong  acids  and  also  in 
alkalies,  so  that  basic  as  well  as  acidic  properties  must  be  ascribed 
to  it  (platinic  acid).  Salts  of  this  acid  are  moreover  formed  when 
platinum  is  fused  with  alkalies.  Platinous  chloride,  PtCl2,  is 
produced  by  heating  chlorplatinic  acid  to  200°  and  in  small  amount, 
also,  when  the  solution  of  this  acid  is  strongly  concentrated.  It  is 
a  green  powder,  insoluble  in  water.  With  the  alkali  chlorides  it 
gives  soluble  double  salts,  such  as  PtCl2-2NaCL  Double  cyanides 
of  platinum  with  many  metals  are  also  known,  e.g.  K2Pt(CN)4* 
3H2O,  BaPt(CN)4-4H2O,  etc.  The  latter  has  come  into  prominence 
because  of  its  ability  to  make  Rontgen  rays  visible.  All  these 
double  salts  are  noted  for  their  beautiful  colors  and  strong  dichroism. 


METAL-AMMONIA    COMPOUNDS.      WERNER'S    EXTENSIONS    OF 
THE  NOTION  OF  VALENCE. 

317.  Several  metals,  notably  those  of  the  eighth  group  of  the  periodic 
system,  are  capable  of  forming  complex  compounds  with  ammonia  and  acid 
radicals.  Such  compounds  have  long  been  known,  some  having  been  pre- 
pared by  the  old  master,  BERZELIUS.  The  study  of  these  substances  occupied 
various  investigators  of  the  nineteenth  century,  especially  JORGENSEN. 
In  recent  years,  however,  this  field  has  been  explored  and  greatly  extended 
by  the  investigations  of  WERNER  and  his  pupils,  so  that  at  the  present  time 
over  1700  compounds  of  the  general  type  MXp(Am)q,  are  already  known, 
M  being  a  metal  atom,  X  an  acid  radical  and  Am  ammonia  or  an  organic 
base  (or  even  water).  Chief  credit  is  also  due  WERNER  for  having  taken  up 
the  theoretical  study  of  the  relationships  between  these  compounds  and,  as 
a  result,  generalizations  of  considerable  importance  for  the  structure  of 
inorganic  compounds,  especially  the  complex  salts,  have  been  established. 
The  whole  subject  deserves  a  little  attention  at  this  point. 

Concerning  the  methods  of  preparing  these  metal-ammonia  compounds 
very  little  of  a  general  nature  can  yet  be  stated.  It  is  readily  appreciated 
that  the  preparation  of  such  a  large  number  of  complex  compounds  calls  for 
the  most  diversified  synthetical  methods. 


317.]  METAL-AMMONIA  COMPOUNDS.  487 

In  order  to  acquire  an  insight  into  the  nature  of  the  compounds  concerned 
we  may  first  examine  the  trinitrito  triammine  cobalt,  Co(NH3)3(NO2)3, 
which  can  be  obtained  by  mixing  cold  solutions  of  cobalt  chloride,  ammo- 
nium chloride  and  sodium  nitrite  and  treating  the  mixture  with  a  current  of 
air.  The  compound  then  separates  out  as  a  difficultly  soluble  crystalline 
powder.  It  can  be  recrystallized  from  hot  water  containing  a  little  acetic 
acid,  without  liberating  nitrous  acid,  and  is  not  attacked  by  dilute  mineral 
acids  in  the  cold.  The  electrical  conductance  of  its  aqueous  solution  is 
approximately  zero.  Evidently,  therefore,  the  substance  lacks  the  ordinary 
properties  of  a  nitrite;  the  NO2-groups  must  be  joined  to  the  molecule  differ- 
ently than  in  the  nitrites.  The  ammonia,  too,  is  otherwise  combined  than 
in  the  ammonium  salts.  This  follows  at  once  from  the  fact  that  the  com- 
pound is  a  non-electrolyte:  furthermore,  from  the  fact  that  the  action  of 
even  concentrated  acids  is  insufficient  to  split  off  ammonia.  The  NH3-groups 
are  not  held  by  the  acid  radicals  present  in  the  molecule,  for  by  energetic 
reactions  it  is  possible  to  substitute  other  acid  radicals  for  these  without 
liberating  the  ammonia  molecules.  The  supposition  of  an  active  participa- 
tion of  the  acid  radicals  in  the  linkage  of  the  NH3-molecules  is  thus  excluded. 
In  deciding  how  the  ammonias  are  connected  to  the  molecule  it  is  significant 
that  the  ammonia  molecules  can  be  replaced  successively  by  other  molecules. 
This  shows  that  each  ammonia  molecule  must  be  linked  independently  of 
the  others  in  the  complex  molecule.  About  the  only  satisfactory  explana- 
tion is  that  the  NH3-molecules  are  attached  directly  to  the  metal  atom. 
WERNER,  however,  makes  the  same  assumption  for  the  acid  radicals  in  order 
to  explain  their  abnormal  behavior.  This  theory,  whereby  the  special 
properties  of  the  groups  in  compounds  of  the  type  MXp(NH3)q  are  explained 
on  the  assumption  that  these  groups  are  in  direct  combination  with  the 
metal,  has  proved  to  be  of  great  importance  for  the  classification  of  these 
compounds.  The  acceptance  of  this  principle,  however,  necessitates  an 
extension  of  the  present  notion  of  valence.  Cobalt,  for  example,  is  at  most 
trivalent  in  its  salts  and  oxides;  but,  if  we  assume  a  direct  linkage  to  the 
metal  of  the  three  NH3-  and  the  three  NO2-groups,  the  cobalt  must  have  a 
valence  of  six.  These  new-appearing  affinities  of  the  metal  atom  cannot 
offhand  be  classed  with  the  ordinary  valence  bonds.  For  this  reason  it 
seems  appropriate  to  assign  a  special  name  to  them.  In  order  to  distinguish 
them  from  the  ordinary  valences,  which  are  termed  "principal"  or  "primary" 
valences,  they  are  called  "subordinate,"  or  "secondary,"  valences. 

Compounds  of  the  type  of  the  trinitrito  triammine  cobalt,  that  is,  of 
the  general  formula  MX3(NH3)3  have  the  property  of  combining  with  more 
ammonia  molecules  still.  Thus  there  are  compounds  of  the  types  MX3(NH3)4, 
MX3(NH3)5,  and  MX3(NH3)g.  The  addition  of  ammonia  gives  rise  to  a 
peculiar  change  in  the  function  of  the  acid  radicals,  because  for  every  addi- 
tional NH3-molecule  that  is  taken  on,  one  of  the  acid  radicals  enters  the 
ionizable  condition.  If  we  compare  the  molecular  conductivities  in 
or  -gfa  normal  solution  at  250°,  to  wit: 


488  INORGANIC  CHEMISTRY.  [§§  317- 


Co(NH3)6Xs        Co(NH3)5X3 
a       402  245  117 

with  those  of  the  salts: 


MgCl,  NaCl  - 

a      370  249  125 

ore  find  that  the  first  of  the  above  complex  salts  must  be  a  quaternary,  the 
second  a  ternary,  the  third  a  binary,  electrolyte,  and  the  fourth  a  non-elec- 
trolyte. 

The  chemical  properties  of  these  compounds  accord  perfectly  with  this 
theory.  In  the  compound,  Co(NH3)4Cl3,  formerly  called  "  praseo-cobalt 
chloride,"  only  one  chlorine  can  be  directly  precipitated  by  silver  nitrate; 
in  Co(NH3)5Cl3  two  can  be  so  precipitated,  while  in  Co(NH3)6Cl3,  hexammine 
cobaltic  chloride,  all  three  chlorine  atoms  appear  to  be  ionized  in  aqueous 
solution,  since  all  three  react  at  once  with  a  silver  salt  solution.  This  is 
explained  on  the  assumption  that  the  added  ammonia  molecules  displace 
the  acid  radicals  from  their  immediate  connection  with  the  metal  atom  and 
themselves  enter  into  this  direct  union  with  the  metal.  In  order  to  distin- 
guish the  ammonia  molecules  which  are  linked  up  in  this  way  WERNER 
applies  to  them  the  term  ammine  and  devises  the  following  formulae  and  names 
to  express  the  constitution  of  the  above-named  compounds: 


M(NH3)3X3  Tri-acido  triammine  compounds 
M(NH3)4X3  Di-acido  tetrammine  salts  [M(NH3)4X2]X, 
M(NH3)6X3  Acido  pentammine  saltt  [M(NH3)5X]X2, 
M(NH3)6X3  Hexammine  salts  [M(NH3)6]X3, 

the  atoms  or  groups  within  brackets  being  regarded  as  in  direct  union 
with  the  metal  atom.  The  ionizable  groups  X  outside  of  the  brackets  are 
in  indirect  union  with  the  complex  and,  according  to  WERNER, 
are  not  joined  to  a  definite  elemental  atom. 

If  we  compare  the  composition  of  the  various  metal-ammonias,  we  find 
that  the  formation  of  complexes  is  not  without  its  limitations,  but  that  after 
the  addition  of  a  definite  number  of  NH3-molecules  it  comes  to  an  end. 
Particularly  sharp  is  the  limit  in  respect  to  the  number  of  groups  which  can 
unite  directly  with  an  atom  serving  as  the  center  of  a  complex  radical.  It 
is  a  striking  fact  that  this  limit  is  the  same  for  a  good  many  elements  and 
is  commonly  six  (6).  It  seems  most  likely  that  this  limiting  number  is 
characteristic  of  the  elemental  atom  and  of  considerable  importance.  WER- 
NER calls  it  the  coordination  number  and  defines  it  as  the  maximum  number 
of  individual  groups  that  can  be  directly  united  to  an  elemental  atom.  As 
already  stated,  the  coordination  number  for  most  elements  is  6;  for  a  few 
elements,  boron,  carbon,  and  nitrogen,  it  is  4. 


318.]  METAL-AMMONIA  COMPOUNDS.  489 

318.  The  constitution  deduced  for  the  metal-ammonia  compounds  en- 
ables us  to  explain  in  a  very  simple  way  a  series  of  isomerism  phenomena 
that  are  characteristic  of  these  compounds.  Two  isomeric  compounds  are 
known  of  the  formula  Pt(NH3)4-SO4(OH)2.  One  behaves  as  a  strong  base, 
absorbing  carbon  dioxide  from  the  air  like  caustic  alkalies  and  precipitating 
metallic  oxides  from  their  salts,  but  gives  no  reaction  of  the  sulphate  ion 
in  aqueous  solution,  e.g.,  no  precipitate  of  barium  sulphate  with  barium 
salts.  The  second  compound  is  a  perfectly  neutral  salt  and  acts  as  a  normal 
sulphate,  giving  a  barium  sulphate  precipitate  at  once  with  barium  salts. 
We  conclude  from  these  reactions  that  the  first  compound  gives  only  OH- 
ions ;  the  second,  on  the  contrary,  only  SO4-ions.  The  cause  of  the  isomerism 
is  explained  in  the  following  coordination  formulae; 

[o2S<^>Pt(NH3)4~J  (OH),        and         [~HO/Pt(NH3)<] S(V 

Sulphate  tetrammine  plato  Dihydroxylo  tetrammine 

hydroxide  plate"  sulphate 


The  researches  on  metal-ammonia  compounds  have  been  of  great  assist- 
ance in  clearing  up  the  structure  of  many  complex  salts,  since  the  latter 
can  be  prepared  by  the  gradual  replacement  of  NH3-groups  by  other  mole- 
cules. An  illuminating  example  of  this  is  potassium  cobaltic  nitrite  (§  310), 
whose  relation  to  hexammine  cobaltic  chloride  is  shown  by  the  following 
series  of  ammonia  compounds: 

1.  [CoCNHP      -     2. 


. 

Nitrito  pentammine 
V  cobaltic  chloride 

[    (N02)s"| 

C°(NH3)3J 


Hoxammine  cobaltic  Nitrito  pentammine  Dinitrito  tetrammine 

chloride  V  cobaltic  chloride  cobaltic  chloride 


ro'ic  -»  «     ro'ic 

(KB*!.  CO(NH3)K> 

Trinitrito  triammine  Potassium  tetranitrito  (unknown) 

cobalt  diammine  cobaltite 

7.  [Co(N02)6]K3. 

Potass  urn  cobaltic  nitrite 

In  all  these  compounds  there  is  no  ionizable  NO2-radical.  The  existence 
of  such  transition  series  clearly  brings  out  the  constitutional  relationships 
which  must  exist  between  the  metal-ammonias  and  the  complex  salts.  The 
formulae  for  these  are  so  constructed  that  the  groups  which  do  not  respond 
to  analytical  tests,  as  well  as  the  NH3-molecules  displaced  by  them,  are  in 
direct  union  with  the  central  metal  atom. 

The  progressive  substitution  of  NH3  by  a  negative  group  or  element 
results  in  the  gradual  decrease  of  the  valence  of  the  cation.     While  this  is 


490 


INORGANIC  CHEMISTRY. 


[§  318. 


three  in  No.  1,  it  has  become  zero  in  No.  4,  i.e.,  the  molecule  has  become  a 
non-conductor.  If  the  substitution  is  carried  further,  as  in  Nos.  5,  6,  and  7, 
the  complex  in  brackets  assumes  the  anion  role  and  its  negative  charge 


1234567 
FIG.  77. 

increases  from  one  to  three.  If  the  molecular  conductivity  for  a  given 
concentration  is  plotted  on  the  ordinate  axis  in  a  diagram,  we  obtain  for 
the  series  of  compounds  a  curve,  as  is  shown  in  Fig.  77.  Numerous  analo- 
gous transitions  of  metal-ammonia  compounds  to  complex  salts  have  been 
discovered,  e.g.: 


[Pt(NH3)6]Cl4 


[PtCl6]K2;          [Pt(NH3)JCl2 


[PtClJK2;  etc. 


Furthermore,  relationships  have  been  found  to  exist  between  metal- 
ammonias  and  hydrates.  It  is  a  significant  fact  that  for  a  large  number 
of  salts  of  metals  which  give  hexammine  salts  with  ammonia,  hydrates  with 
six  molecules  of  water  predominate.  This  is  the  case  with  the  salts  of 
cobalt,  chromium,  nickel,  etc.  Moreover,  JORGENSEN  observed  that  com- 
pounds are  derived  from  the  hexammine  salts  by  the  exchange  of  one  or 
two  NH3-molecules  for  water,  which  compounds  correspond  in  behavior 
to  the  hexammine  salts.  This  follows  from  the  fact  that  the  molecular 
conductivity  in  aqueous  solution  is  but  little  affected  by  this  exchange,  e.g.: 


[Co(NH3)6]Br3, 
a  402 


390 


380 


It  is  also  confirmed  by  the  fact  that  with  the  removal  of  each  water  mole- 
cule an  acid  radical  loses  its -ionizing  property,  just  as  was  the  case  for  the  | 
removal  of  each  ammonia  molecule  from  the  hexammine  salts: 


METAL-AMMONIA   COMPOUNDS.  491 

,(H20) 


"      KUJ* 

Chloro  pentammine 


H30. 


Aquo  pentammino  Chloro  pentammine 

chromic  chloride  chromic  chloride 


The  addition  of  water  has  the  opposite  effect.  The  trichloro  triammine 
cobalt,  for  example,  is  known  to  form  hydrates  with  1,  2,  and  3  mols.  water. 
Since  all  of  the  three  chlorine  atoms  in  the  trihydrate  have  an  ionizing  charac- 
ter, while  in  the  diyhdrate  only  two  have  this  property,  in  the  monohydrate 
one  and  in  the  anhydrous  compound  none,  we  are  justified  in  ascribing  to 
them  in  the  following  constitution: 


r  Cl2  i     r  cl  i 

[Cl3Co(NH3)3],       Co(H20)     Cl,       Co(H20)2   C12,    and 
L     (NH3)3-I          L     (NH3)3J 

A  nearly  complete  transition  series  from  a  metal-ammonia  to  a  hydrate 
(hydrous  salt)  is  to  be  found  in  the  case  of  chromium: 

r  (H2o)  i 

3'        Cr(NH3)5JCl3' 


Lacking  Blue  hexahydrate 


AH  these  compounds  contain  three  chlorine  atoms  in  ionizable  linkage; 
accordingly  silver  nitrate  at  once  precipitates  all  the  chlorine  as  silver 
chloride,  even  from  freshly  prepared  solutions  of  the  salts.  The  function  of 
water  in  these  compounds  corresponds  perfectly  to  that  of  the  water  in  the 
aquo-salts  of  the  cobalt  ammonias,  i.e.,  with  the  release  of  each  water  mole- 
cule a  chlorine  atom  sacrifices  its  ionizing  ability. 

Two  hexahydrates  of  chromic  chloride,  CrCl3-6H2O,  are  known,  one  blue 
and  the  other  green.  The  blue  one  contains  three  ionizable  Cl-atoms,  and 
according  to  its  general  behavior  should  be  regarded  as  [Cr(OH2)6]Cl3,  hexaquo 
chromic  chloride.  Upon  the  loss  of  two  water  molecules  it  goes  over  into 
dichloro  tetraquo  chromic  chloride  [Cl2Cr(OH2)4]Cl,  which  contains  but  one 
ionizable  Cl-atom. 

If  this  dichloro  tetraquo  chloride  is  crystallized  out  of  water,  it  adds 
on  two  molecules  of  the  water  and  goes  over  into  the  green  hydrate, 
[Cl2Cr(OH2)JCl  +  2H2O,  which  is  thus  isomeric  with  the  blue  hydrate.  The 
green  hydrate,  moreover,  contains  only  one  chlorine  atom  that  can  be  pre- 
cipitated directly  with  silver  nitrate.  It  should  also  be  mentioned  that 
the  green  hydrate  can  be  transformed  into  the  blue  one  and  the  water  mole- 
cules, that  condition  the  isomerism,  shifted.  Since  the  isomerism  must 


492 


INORGANIC  CHEMISTRY. 


[§318, 


depend  on  the  difference  in  the  manner  of  attachment  of  the  water,  WERNER 
has  called  it  hydrate  isomerism. 

Various  cases  of  isomerism  that  are  met  with  in  the  metal-ammonia 
compounds  are  best  explained  as  cases  of  stereo   isomerism.     Two  series 

of  compounds  are  known  of  the  general  formula,  Pt     v3  2J   tnev  are  called 
platin  diammine  compounds  and  have  been  known  since  1870.     Later  in- 


--r 


B 

FIG.  78. 


FIG.  79. 


vestigations  have  shown  that  this  kina  of  isomerism  occurs  only  with  com- 
pounds of  the  type  M    2.     The  assumption  can  therefore  be  made  that  the 

X54 

six  groups  are  arranged  at  the  apexes  of  an  octahedron  around  the  central 

atom.     Only  with  the  type  above  described  are  two  isomers  possible,  as 

^ 

the  accompanying  figures  (Figs.  78  and  79)  show.      Compounds  of  the  M    * 

r> 

type  cannot,  therefore,  have  isomers  and,  as  a  matter  of  fact,  no  one  has  yet 
been  able  to  prepare  such  isomers. 

WERNER  has  also  discovered  metal-ammonias  of  more  than  one  nucleus, 
i.e.,  where  the  molecules  contain  several  metal  atoms  joined  to  one  another 
in  such  a  way  as  to  undergo  no  separation  either  by  ionization  or  by  spon- 
taneous hydrolysis;  but  we  cannot  dwell  longer  upon  the  subject  her*. 


INDEX. 


The  principal  references  are  in  italics. 


Abraum  salts  of  Stassfurt,  65,  322, 

374,  441 

Absolute  weight  of  atoms,  49 
Absorption,  70 
Absorption  spectra,  391 
Accumulator,  421 
Acetylene,  249 
Acidimetry,  350 
Acids,  39,  99 

strength  of,  99 
Acid  chlorides,  198 
Acids,  dibasic,  131 
Acids,  monobasic,  83 
Acid  salts,  131 
Acid  sodium  carbonate,  322 
Actinium,  399 
Additive  properties,  348 
Affinity,  156 
Agglutinants,  269 
Air,  165 
Alabaster,  376 
Alchemists,  371 
Aldebaranium,  446 
Algaroth  powder,  234 
Alkalimetry,  350 
Allotropism,  53,  104,  148,  201,  222, 

241,  261,  274 
Alloys,  232,  237,  276,  355,  361,  409, 

437,  471,  479,  482,  483,  484 
Aludels,  410 
Alumina,  437 
Aluminates,  438 
Aluminium,  436 

amalgam,  437 

bronze,  437 

chloride,  438 

hydroxide,'  438 

oxide,  437 

silicate,  440 

sulphate,  439 
Alums,  43& 
Alum  stone,  439 
Alunite,  439 
Amalgams,  412 
Amalgam,  aluminium,  437 


Amalgam,  ammonium,  175  j 

sodium,  313 

tin,  276 

Amblygonite,  310 
Amethyst,  265 
Amides,  .198 

Amidophosphoric  acid,  221 
Amido-sulphonic  acid,  198 
Ammine,  488 
Ammonia,  174,  311 

compounds  of  the  metals  of  the 
eighth  group,  486 

-soda  process,  320 
Ammonium,  175,  331 

amalgam,  175 

carbonate,  334 

chloride,  333 

hydroxide,  176 

magnesium  phosphate,  216 

metavanadate,  448 

nitrate,  334,  379 

phosphate,  334 

salts  in  analysis,  375 

sulphate,  334 

sulphide,  335 

Amorphous  substance,  385 
Amphoteric  compounds,  278,  284 
Amphoteric  reaction,  319 
Analysis,  hydrogen  sulphide  in,  117 
Analysis,  volumetric.     See  Volumet- 
ric analysis. 
Anatase,  446 
ANAXAJGORAS,  28 
Anglesite,  286 
Anhydride,  acid,  83 

mixeti,  186,  195 
Anhydfrite,  381 
Anions;  96 
Anode,  96 
Antichlor,  132,  317 
Antimonic  acid,  236 
Antimony,  231 

butter,  233 

pentachloride,  234 

pentasulphide,  236 

493 


494 


INDEX. 


Antimony,  pentoxide,  236 

tetroxide,  236 

trichloride,  233 

trioxide,  234 

trisulphide,  236 
Antimonyl,  235 
Antiseptic,  416 
Anglesite,  286 

Angstrom,  392 
Apatite,  199,  376 
Aqua  regia,  198 
Aragonite,  376,  382 
Argentite,  359 
Argon,  170,  300,  309 
Argyrodite,  286 
ARISTOTLE,  371 
ARRHENIUS,  95,  101 
Arsenic,  221 

acid,  229 

meal,  226 

oxide,  228 

sulphides,  229,  268 

sulpho-salts,  230 

trichloride,  226 

white,  224,  226 
Arseni-tungstates,  457 
Arsenious  acid,  228 

oxide,  226 
Arsenolite,  221 
Arsenopyrite,  221 
Arsine,  223 
Asbestos,  374 
Association,  82,  185 
Atmosphefe,  165 
Atomic  heat,  289 

theory,  27 

volume,  297 

weights,  29     v 

Determination  of,  44,  47,  49,  288, 

305 

Table,  of,  Inside  back  cover 
Atoms,  28,  49 
Augite,  374 
Auric  compounds,  370 
Aurous  compounds,  369 
Aurum  musivum,  281 
AVOGADRO'S  hypothesis,  44,  60,  293 
Azides,  179 
Azurite,  353 

Bacteria,  194 
BAEYER,  144 
BAKER,  185,  333,  414 
BALARD,  65,  84 
BALMER,  392,  393 
Barite,  387 
Barium,  387 


Barium  peroxide,  53,  387 

Baryta-water,  387 

Bases,  39,  102 

BASILIUS  VALENTINUS,  231 

BAUME  degrees,  142 

Bauxite,  436,  438 

BECKER,  129 

BECQUEREL,  397 

BeU  metal,  276 

BERGMANN,  99 

BERNTHSEN,  133 

BERTHELOT,  50,  152,  157,  182,  255 

BERTHELOT'S  principle,  157 

BERTHELOT-MAHLER        calorimetric 

bomb,  153 
Beryl,  372 
BeryUium,  303,  372 
BERZELIUS,  23,  29,  30,  148,  280,  293, 

308,  418,  486 
BESSEMER  process,  469 
Bimolecular  reactions,  76,  206,  255 

BlRKELAND,   191 

Bismuth,  236 

chloride,  237 

glance,  237 

hydroxides,  238 

oxides,  238 

oxychloride,  237 

subnitrate,  239 

sulphate,  239 

sulphides,  239 
Black  ash,  319 
Blanc  fixe,  388 
Blast  furnace,  463 
Bleaching,  36,  56,  126 
Bleaching-powder,  85,  380^ 
Blende,  407  ^^^ 

Blooms,  471 

BODENSTEIN,  15 

Boiling-point,  elevation  of,  61,  272 

BOLTWOOD,  404 

BONE,  249 

Boneblack,  244 

Boracite,  432 

Borax,  432,  435 

Boric  acid,  433 

Boron,  432 

BOUSSINGAULT,  166 

BOYLE,  371,  395 

BOYLE'S  law,  47,  58,  160,  295 

BRAND,  199 

Brass,  355 

Bricks,  440 

Brimstone,  103 

Bromine,  65 

oxygen  compounds  of,  90 
Bronze,  phosphor,  202,  276 

silicon,  276 


INDEX. 


495 


Bronzes,  276 

Brookite,  446 

Brownian  movement,  49 

BROWNING,  J.,  spectroscope,  389 

BRUHL,  55 

BUNSEN,  329,  389,  391,  398 

burner,  258 

ceU,  422 
Burette,  147 

Cadmium,  410 
Caesium,  329 
CAIN,  249 
Calamine,  407 
Calcining,  353,  463 
Calcite,  376,  382 
Calcium,  376 

carbide,  171,  249 

carbonate,  382 

chloride,  321,  340,  379 

cyanimide,  175 

fluoride,  380 

hydroxide,  378 

hypochlorite,  380 

manganite,  460 

nitrate,  382 

oxide,  377 

peroxide,  379 

phosphates,  382 

plumbate,  285 

silicate,  384 

sulphate,  338,  381 

sulphide,  291,  319,  380 
Calomel,  413 

Caloric  effect,  152,  336,  357,  437 
Calorimeter,  153 
Calx,  165 
Canal  rays,  400 
Carat,  368 

Carbides,  247,  249,  457 
Carbon,  241 

amorphous,  244 

dioxide,  252 

monoxide,  249 

oxysulphide,  251 

tetrafluoride,  247 
Carbonado,  241 
Carbonic  acid,  253 
Carborundum,  247 
Carburetting,  250 
Carnallite,  322,  329,  374,  441 
CARO'S  acid,  144 
Cassiopeium,  446 
Cassiterite,  274 
Cast  iron,  466 
CASTNER,  process,  314 
Catalysis,  34,  55,    73,    129,  134,  137, 
181,  194,  233,  334,  380 


Cations,  96 
Cats-eye,  455 
Caustic  potash,  323 

soda,  313 

CAVENDISH,  17,  165 
Celestite,  386 
CELSIUS,  20 
Cement,  378 
Cementite,  466 
Ceramics,  440 
Cerargyrite,  359 
Cerite,  443 
Cerium,  445 
Cerussite,  281 
Chalcocite,  353 
Chalcopyrite,  103,  353 
Chalk,  376 

Chamseleon  minerale,  461 
Chamber  acid,  136,  319 

crystals,  195 
Charcoal,  245 
CHARPY,  466 
Chemical  affinity,  36,  156 

operations,  5 
Chemistry,  field  of,  3 
Chili  saltpetre,  69,  189,  311,  318,  326 
Chinese  white,  409 
Chloramine,  177 
Chlor-auric  acid,  370 
Chlorazide,  179 
Chloric  acid,  88 
Chloride  of  lime,  380 
Chlorination  process  (gold),  368 
Chlorine,  33 

hydrate,  37 

oxygen  compounds  of,  83 
Chlorous  acid,  88 
Chlorosulphonic  acid,  142 
Chlorplatinic  acid,  485 
Chromates,  452 
Chrome  alum,  452 

green,  450 

yellow,  454 

Chrom-sulphuric  acids,  451 
Chromic  anhydride,  452 

chloride,  451 

oxide,  450 

sulphate,  451 
Chromite,  449,  450 
Chromium,  449 

in  steel,  471 

oxychloride,  455 
Chromous  compounds,  450' 
Chromyl  chloride,  455 
Chryaoberyl,  372 
Cinnabar,  103,  410 
CLARKE,  386 
Clay,  436,  440 


496 


INDEX. 


Cleveite,  172 
Coal,  245 
Coarse  metal,  354 
Cobalt,  477 

compounds  (ammonia),  487 
Cobaltite,  221,  477 
COEHN,  129 
COHEN,  425 
Coke,  245 
Colcothar,  474 

Cold-hot  tube  of  DEVILLE,  128 
Colemanite,  432,  435 
Collargol,  269 

Colloids,  266,  268,  365,  438,  474 
Columbium,  449 
Combustion,  11,  165 
Components,  110 
Compound,  25 
Condenser,  6 

Conductivity,  molecular,  97 
Conservation  of  matter,  16 

energy,  153 

Constancy  of  natural  phenomena,  3 
Contact  action.     See  Catalysis. 

process  (sulphuric  acid),  137 
Converter,  469 
Coordination  number,  488 
Copper,  353 

ammonia  compounds,  359 

compounds,  355 

(See  "Cupric"  and  "Cuprous.") 
Copperas,  473 

Corrosive  sublimate,  414,  415 
Corundum,  436 
Counter-current  principle,  19 
COURTOIS,  69 
Covering  power,  287 
CREDNER,  168 
Crocoite,  281,  449 
CROOKES,  442 

Crucible  process  (steel),  470 
Cryohydric  point,  339,  341,  379 
Cryohydrate,  339,  343 
Cryolite,  311,  436 
Cryoscopic  method,  65 
Crystallization,  6 
Crystalloids,  268 
Crystals,  mixed,  386,  468 
Cup  and  cone  apparatus,  463 
Cupellation,  360,  368,  369 
Cupric  arsenite,  359 

bromide,  358 

carbonate,  359 

chloride,  358 

hydroxide,  358 

iodide,  358 

nitrate,  359 

oxide,  358 


Cupric  sulphate,  358 

sulphide,  359 
Cuprite,  353 
Cuprous  chloride,  356 

cyanide,  357 

iodide,  356 

oxide,  355 

sulphate,  357 
CURIE,  397,  398 
CURTIUS,  178 
Cyanide  process,  360,  368 
Cyanogen,  255 

DALTON,  27,  43,  293,  397 
DANIELL  cell,  420 
DAVY,  241,  258,  311,  322,  418 
DEACON  process  (chlorine),  34 
DEBIERNE,  398 
Decantation,  5 
Decay  of  elements,  402,  404 
Definite  proportions,  law  of,  27 
.Degree  of  freedom,  111 
Deliquescence,  318 
DEMOCRITES,  28 
Dephlogisticated  air,  166 
Desiccators,  142 
Detonating-gas,  15,  22 
DEVILLE,  STE.  CLAIRE,  73,  127 
DEWAR,  14,  172 
Dialysis,  266,  269 
Diamide,  177 
Diamond,  241 
Didymium,  445 
Dilatometer,  108 
Disinfection,  36,  416,  473 
-  — Pisperse,  phase,  etc.,  272 

Dissociation,  70,  72,  94,  121, 124, 140, 

148,  185,  193,  197,  210,  227,  233, 

259,  277,  286,  333,  355,  370,  383, 

414,  474 
electrolytic,  94,  131,  214,  219,  254, 

268,  313,  332,  346,  357,  363,  375, 

409,  415,  417,  418,  472,  475,  478, 

489. 
hydrolytic,  100,  207,  269,  277,  279, 

352,  362,  376,  414,  433,  435,  439, 

442,  447,  474 
tension,  430 
Distillation,  £,19 

fractional,  6 
Distribution  law  of  BERTHELOT,  325. 

332,  360 

Dithionic  acid,  145 
Divariant  system,  114 
DIXON,  255 
DOEBEREINER,  212 
Dolomite,  374 
Double  decomposition,  41 


INDEX. 


497 


Double  salts,  278,  286,  360,  440,  486 

DUBOIS,  168 

DULONG  and  PETIT,  law  of,  289,  407 

DUMAS,  23,  24,  166 

Dysprosium,  443 

Earthenware,  440 
Earth's  crust,  composition  of,  8 
Eau  de  Javelle,  85 
Ebullioscopic  method,  65 
Efflorescence,  318,  321 
EINSTEIN,  272 
Eka-aluminium,  441 
-boron,  445 
-silicon,  273,  307 
Electrical  endosmose,  271 
Electric  furnace,  243,  249,  262,  265, 

377,  457,  465,  470 
Electrochemical  series,  428 
Electrochemistry,  418 
Electrolysis,  23,  41,  79,  133,  143,  177, 

310,  311,  313,  322,  325,  354,  360, 

361,  362,  368,  386,  387,  430,  436,' 

447,  454,  457,  485 
Electrolytic  separation  of  the  metals, 

430 

Electromotive  force,  159,  419,  426 
Electrons,  396 
Electrotyping,  355 
Elements,  7,  296,  371 
Emanation,  401 
Emerald,  372 
Emery,  436 
EMPEDOCLES,  28,  156 
Emulsion,  272 
Endothermic  compounds,   153,   161, 

179,  180,  182,  184,  187,  223,  249 
Energy,  free,  bound,  and  total,  157, 

158 

Epsom  salt,  376 
Equations,  31 
Equilibrium,  73,  86,  98,  118,  128,  138, 

157,  159,  160,  177,  186,  216,  226, 

250,  259,  325,  334,  337,  355,  356, 

375,  379,  383,  418,  465,  467,  472 
heterogenous,  109,  383 
state  of,  77 

Equivalent  weights,  293,  305 
Erbia,  445 
ERDMANN,  170 

Eudiometric  analysis  of  air,  166 
Europium,  443 
Eutectic  point,  339,  343,  468 
Euxenite,  443 
Exothermic  reactions,  153,  155,  158, 

161,  357 

Explosion  wave,  255 
EYDE,  191 


FARADAY,  law  of,  159,  397 
Feldspar,  436 
Ferric  compounds,  474 
Ferrite,  466 
Ferrocyanic  acid,  476 
Ferrous  compounds,  473 
Feuerluft,  166 
Filtration,  5 
Fire  damp,  248 
Flame,  256 

spectrum,  390 
Flash  light,  374 
Flowers  of  sulphur,  103 
Fluorine,  79  / 

Fluor  spar,  79,  81,  380 
Flux,  353,  463 
Formula,  30 

empirical,  57 

Fractional  distillation,  6        N 
FRASCH  sulphur  process,  103 
FRAUNHOFER  lines,  394 
Freezing  mixture,  379 

-point,  depression  of,  61,  94,  105, 

121,  272,  339 

Freezing-point  curves.     See  Melting- 
point  curves. 
FR:MY,  308 
Fuels,  245 
Fumaroles,  433 

Furnace,  electric.     See  Electric  Fur- 
nace. 

reverberatory,  360 , 

Gadolinite,  443 
Galenite,  103,  281 
Gallium,  441 

Galvanic  cells,  418  ^ 

Galvanized  iron,  409,  424) 
Garnierite,  479 
Gas  carbon,  245 
GATTERMANN,  71 
GAY-LUSSAC,  69,  87,  192,  311) 

law  of,  43,  47,  58 

tower,  136,  196 
GEBER,  198 
Germanium,  273,  307 
German  silver,  355,  479  ^ 
Gersdorffite,  479 
GIBBS,  109 

Glacial  phosphoric  acid,  217 
Glass,  384 

etching,  82,  263 
GLAUBER,  156 
GLAUBER'S    salt.     See    Sodium    suit 

phate. 

GLOVER  tower,  135,  196 
Glucinum,  372 
Gold,  269,  367 


498 


INDEX. 


Gold  compounds.     See  Aurous  com- 
pounds and  Auric  compounds. 

Gold-plating,  368 

GOLDSCHMIDT,  reduction  method,  437 
449,  456,  458 

GOLDSTEIN,  400 

Graduation  salt  process,  315 

GRAHAM,  266,  268 

Gram  equivalent,  146 
molecule,  32 

Graphite,  244 

Graphitic  acid,  244 

GRIESHEIM  process,  313,  325 

GROVE'S  gas  battery,  425 

GUIGNET'S  green,  450 

Gun-metal,  276 

Gunpowder,  26,  327 

GUTZEIT'S  test  (arsenic),  222,  224 

GUYE,  294 

GYPSUM,  103,  376,  381 

Halite,  311 

HALL  process  (aluminium),  436 
Hammer  scale,  12 
HAMPSON,  10,  170 
Hardness  of  water,  383 
HARGREAVE'S  method,  317 
Hartshorn,  salt  of,  334 
Hausmannite,  458 
Heat,  atomic,  290 

molecular,  173,  291 

of  dilution,  153 
formation,  152,  155 
neutralization,  153,  350 
solution,  153,  161,  336 

specific,  290 
Heavy  spar,  387 
Helium,  172,  300,  394,  400,  402 
HELMHOLTZ,  158 
Hematite,  463 

HENRY'S  law,  11,  39,  95, 133,  332,  434 
Hepar  sulphuris,  328 
HEROULT,  436 

furnace,  470 
HESS,  law  of,  153 
Hexammine  cobalt  salts,  489 
Hittorf,  202 
HOFMANN,  24 
Holmium,  443 
Hopper  crystal  of  salt,  317 
Hornblende,  374 
Horn  silver,  359,  363 
Hiibnerite,  456 
Hydrate  isomerism,  192 
Hydrates,  490 
Hydraulic  mining,  368 
Hydrazine,  177 
Hydrazoic  acid,  178 


Hydriodic  acid,  71 
Hydrobromic  acid,  67 
Hydrocarbons,  248 
Hydrochloric  acid,  37 

composition  of,  41 
Hydrofluoboric  acid,  433 
Hydrofluoric  acid,  81 
Hydrofluosilicic  acid,  264 
Hydrogel,  270,  358,  438,  450,  474 
Hydrogen,  13 

antimonide,  232 

arsenide,  223 

bromide,  67 

chloride,  37 

cyanide,  256 

disulphide,  120 

fluoride,  81 

iodide,  71 

peroxide,  53,  64,  367,  455 

persulphide,  120 

phosphide,  204 

selenide,  149 

silicide,  262 

sulphide,  115 

telluride,  150 

trisulphide,  121 
Hydrolysis.  See  Dissociation,  hydro- 

lytio. 

Hydrosol,  270 
Hydroxyl,  143 
Hydroxylamine,  180 

-disulphonic  acid,  198 
Hypo,  132,  317 
Hypochlorous  acid,  84 

oxide,  84 

Hyponitrous  acid,  187 
Hypophosphoric  acid,  218 
Hypophosphorous  acid,  220 
Hyposulphurous  acid,  133 
Hypothesis,  2 

Ice-machine,  175 
Ice  stone,  311,  436 
Illuminating  gas,  476 
Incandescent  electric  light,  456 

§as  light  (WELSBACH),  443 
elible  ink,  366 
Indicators,  350,  352 
Indium,  305,  441 
Induced  radio-activity,  401 
Inversion,  point,  of,  106 
Iodine,  69 
chlorides,  83 

oxygen  compounds  of,  91 
lodometry,  146,  226,  228 
Ionic  equilibrium,  98 

equation,  118,  132,  277,  282,  313, 
356,  408,  422,  453,  461,  475 


INDEX. 


499 


Ionic  equilibrium,  theory  of,  94,  118 

Ionium,  405 

lonization.     See    Dissociation,    elec- 

trolytic. 
Ions,  95,  346,  356,  363,  408,  409,  439, 

451,  453,  459,  485,  487 
valence  of,  123 
Iridium,  484 
Iridosmine,  482,  484 
Iron,  463 

and  carbon  monoxide,  477 
Isomerism,  492 
Isomorphism,  292 

ic  solutions,  62 


Jasper,  265 
JORGENSEN,  486,  490 

Kainite,  322,  374 

Kaolin,  436,  440 

KASSNER,  process,  285 

KAYSER,  392,  393 

Kelp,  69 

Kieserite,  374 

Kinetic  theory,  47,  249 

KIPP  generator,  115 

KIRCHOFF,  329,  389,  391,  398 

KNIETSCH,  137 

KOHLRAUSCH,  19 

KOPP,  291 

KRYPTON,  172,  300 

KUHNE  method,  261 

"Labile,"  108 
LADENBURG,  53 
Lampblack,  245 
Lanthanum,  445 
Lapis  lazuli,  440 
LAVOISIER,  10,  13,  18,  165,  241 
Law  of  AVOGADRO,  44,  47,  60,  293 
BOYLE.     See  BOYLE'S  law. 
the  constancy  of  natural  phe- 

nomena, 3 
onstant    composition    (definite 

proportions),  27 
distribution    (BERTHELOT),  325 

332 

DULONG  and  PETIT,  290,  407 
GAY-LUSSAC,  43,  47,  58 
HENRY.     See  HENRY'S  law. 
HESS,  153 

MlTSCHERLICH,  292 

multiple  proportions,  29 
chemical  mass  action,  75 
NEUMANN,  291 
octaves,  296 
thermoneutrality,  350,  415 


Lead,  281 

carbonate,  286 

chamber  crystals,  195 

chambers,  135 

chloride,  285 

chromate,  454 

glass,  384 

nitrate,  286 

oxides,  283 

peroxide,  294,  422 

persulphate,  286 

sulphate,  286 

sulphide,  287 

tree,  282 

white,  286 

LE  BLANC  soda  process,  142,  319,  327 
LE  CHATELIER'S  rule,  160.  186,  250, 

338 

LECLANCHE  cell,  423 
LECOQ  DE  BOISBAUDRAN,  441 
Lepidolite,  310,  329 
LEUCIPPUS,  28 

LEYDENFROST,  phenomenon,  170 
Lime,  377 

-sulphur  solution,  381 
Limestone,  241,  376 
LINDE,  10,  170 
Liquation,  274 
Liquefaction  of  oxygen,  10 
Litharge,  283 
Lithia  mica,  310,  329 
Lithium,  310 
LOBRY  DE  BRUYN,  178,  272 

LOCKYER,   172 

Lunar  caustic,  366 
LUPKE  cell,  423 
Lutecium,  443 


Magnalium,  437 
Magnesia,  374 

alba,  376 

mixture,  229 

usta,  374 
Magnesite,  374 
Magnesium,  374 

ammonium  phosphate,  376 

boride,  432 

carbonate,  376 

chloride,  375 

hydroxide,  374 

nitride,  174,  374 

oxide,  374 

sulphate,  376 
Magnetite,  463,  474 
Malachite,  353 
Malleable  iron,  466 
MANCHOT,  55 


500 


INDEX. 


Manganese,  458 

dioxide,  33,  460 

in  steel,  471 
Manganic  compounds,  459 

acid,  460 

Manganous  compounds,  459 
Marble,  376 
MARCKWALD,  399 
Marl,  376 
Marsh  gas,  248 
MARSH  test,  225,  233 
Martensite,  466 

Mass  action  law,  75.     See  also  Equi- 
librium. 
Massicot,  283 
MASSON,  299 
Matches,  203 
Matte,  354 
Matter,  3 

unity  of,  308,  395 
Meerschaum,  374 
Mellitic  acid,  246 
Melting-point  curve,   122,  342,  412, 

467 
MENDEL::EFF,  296,  299,  303,  307,  441, 

445 

MENDEL  EEFF'S  table   (periodic  sys- 
tem), 301 
MENSCHING,  232 
Mercuric  ammonium  chloride,  414 

chloride,  414 

cyanide,  268,  414 

halides,  alkali,  415 

iodide,  415 

nitrate,  416 

oxide,  9,  408 

sulphate,  416 

sulphide,  417 

Mercurous  compounds,  413 
Mercury,  410 
Mesothorium,  448 
Metal-ammonia  compounds,  486 
Metalloids,  8 
Metallurgy  of  iron,  463 
Metaphosphoric  acid,  217 
Metaphosphorous  acid,  219 
Metastable  system,  108,  346 
Metastannic  acid,  280 
Methane,  248 
Methyl  orange,  353 
MEYER,  LOTHAR,  41,  296,  300,  308 
MEYER,  VICTOR,  232 
Mica,  436 
MICHELSON,  392 
Microcosmic  salt,  334 
Mineralization  of  organic  matter,  168 
Miner's  safety  lamp,  258 
Minium,  284 


Mispickel,  221 
MITSCHERLICH'S  law,  292 

test,  203 
Mixture,  26,  169 
Mobile   equilibrium,    VAN'T  HOFF'S 

principle  of,  160,  336 
MOHR'S  salt,  473 
MOISSAN,  16,  38,  67,  79,  133,  241, 

244,  247,  446,  449 
Mole,  32 

Molecular  depression,  65 
elevation,  65 
heat,  173,  291 
weight,  45 

determination  of,  47 

by  boiling-point  method,  61. 

227      ...  . 
by  freezing-point  method,  65, 

98,  131,  208,  312 
Molecule,  28,  44,  49 
Molybdenite,  455 
Molybdenum,  455 

trioxide,  455 
Molybdic  acid,  456 
Monazite  sand,  443 
Monobasic  acid,  83 
MORLEY,  29,  212,  294 
Mortar,  378 
Mosaic  gold,  281 
Muriatic  acid,  38 
MUTHMANN,  456 

Nascent  state,  37,  54 

Natural  gas,  248 

Negative  (photography),  365 

Neodymium,  445 

Neon,  172,  300 

NERNST,  104,  184,  282,  418,  425 

glower,  51,  54 
NESSLER'S  solution,  415 
NEUMANN'S  law,  291 
NEWLANDS,  296 
NEWTON'S  metal,  237 
Niccolite,  479 
Nickel,  479 

steel,  471 
Niobium,  449 
Niton,  402 
Nitramide,  199 
Nitric  acid,  189 

oxide,  183 
Nitrides,  164 
Nitrilosulphonates,  198 
Nitrogen,  162,  170 

acid  derivatives  of,  195 

dioxide,  185 

halogen  compounds  of,  179 

hydrogen  compounds  of,  174 


INDEX. 


505 


Touchneedles,  369 

Touchstone,  369 

Tourmaline,  374 

TOWNSEND  process,  314 

Transition  point,  106,  160,  227,  275, 

317,  334,  340,  415,  425 
TRAUBE,  54 
TRAVERS,  172 
Triads,  296 

Triboluminescence,  227 
Tridymite,  265 

Trinitrito  triammine  cobalt,  487 
Triple  point,  113 
Triphylite,  310 
Tuffstone,  378 
Tungsten,  456 

steel,  456,  471 
Twyers,  463 
TYNDALL,  effect,  271 
Type  metal,  232 

Ultramarine,  440 
Ultramicroscope,  272 
Ultraviolet  light,  18,  37,  129,  266 
Unimolecular  reactions,  76,  251,  402 
Unity  of  matter,  395 
Univarient  system,  112 
Uraninite,  397,  457 
Uranium,  177,  397  and  foil.,  457 
URBAIN,  444 

Vacuum  flask,  11,  170 
Valence,  122,  185,  486 
Valence,  maximum,  123 

of  ions,  123 
Vanadinite,  448 
Vanadium,  448 
VAN  DER  STADT,  206,  219 
VAN  DER  WALLS,  49,  109 
VAN  MARUM,  49,  203 
VAN'T  HOFF,  58,  60,  157,  381 
VAN'T   HOFF'S    principle    of   mobile 

equilibrium,  160,  336,  337 
Vapor  pressure  curve,  61,  106,  112 
Varec,  69 
Vein  mining,  367 
Velocity  constant,  75 

of  reaction,  15,  75,  182,  233.     See 

also_Catalysis. 


Vermilion,  417 
Vitriol,  blue,  358 

green,  473 

oil  of,  137 
Vivianite,  199 
VOGEL'S  spectroscope,  389 
Volumetric  analysis,  146,   189,  350, 

460,  462 
Vulcanizing  rubber,  121 

WACKENRODER'S  liquid,  145 
Washing  (precipitates),  5 
Washing-soda,  321 
Water,  17 

composition  of,  22 

gas,  250 

glass,  266,  322,  328' 

natural,  20,  383,  473 

physical  properties  of,  20 

purification  of,  21,  283 
Wavellite,  199 
Welding,  437 
WELDON  process,  460 
WELSBACH,  AUER  VON,  445,  446 

incandescent  gas  light,  250,  257, 

388,  443,  447 
WERNER,  486 
WHITNEY,  347 
WINKLER,  273,  307 
Witherite,  387 
Wolframite,  456 
WOOD'S  metal,  237 
WOULFF  bottle,  18 
Wrought  iron,  466 
Wulfenite,  281,  455 

Xenon,  172,  300 

Ytterbium,  445 
Yttrium,  443 

Zinc,  407 

compounds,  409 

dust,  407 

white,  409 
Zircon,  446 
Zirconia,  447 
Zirconium,  446^ 
ZSIGMONDY,  272 


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